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THE    ELEMENTS 


CHEMICAL    ARITHMETIC 


A  SHORT  SYSTEM  OF  ELEMENTARY 
QUALITATIVE   ANALYSIS. 


BY 

J.   MILNOR   COIT,   PH.D., 

MA8TBB  IN   ST.   PAUL'S   SCHOOL,  CONCORD,   N.H. 


Qt  TH3 

—  '— 


BOSTON: 
D.   C.   HEATH   &  COMPANY. 

1895. 


1 


COPYRIGHT,  MAR.  22,  1886, 
BY  J.  MILNOR  COIT. 


J.  S.  Gushing  &  Co.  — Berwick  &  Smith. 
Boston,  Mass.,  U.S.A. 


PEEFACE. 


THIS  little  manual  is  intended  to  supplement  the 
teaching  of  the  text-books  of  descriptive  chemistry,  and 
to  be  used  as  a  companion  to  them,  by  those  who  desire 
to  make  the  whole  subject  more  practical.  It  is  the 
result  of  the  author's  experience  after  several  years  of 
elementary  science  teaching. 

Part  I.  contains  some  of  the  more  important  rules 
and  principles  of  chemical  arithmetic,  followed  by  a 
series  of  problems,  which  will  not  be  found  to  be  above 
the  comprehension  of  the  average  student  in  the  schools. 
The  matter  relating  to  chemical  theory,  and  the  rules, 
have  been  collected  from  the  best  authorities. 

Part  II.  is  devoted  to  an  elementary  system  of  quali- 
tative analysis,  and  the  best  methods  have  been  adopted. 
This  part  of  the  book  can  be  used  separately,  and  can 
be  taught  together  with  any  good  work  in  descriptive 
chemistry,  such  as  Eliot  &  Storer's,  Shepard's,  RemseSPi, 
or  Avery's  Chemistry.  An  intelligent  student  can,  with 
the  occasional  supervision  of  his  instructor,  work  out  by 
himself  the  reactions  and  the  separations  as  given  in 


iv  PREFACE. 

the  tables.  The  tables  are  those  generally  in  use.  Tests 
are  given  for  the  more  common  metals  and  acids  only, 
and  the  reagents  indicated  are  those  which  almost  any 

school  laboratory  will  afford. 

J.   M.  C. 

ST.  PAUL'S  SCHOOL, 
Concord,  N.H. 


1711 


PAKT  I. 
CHEMICAL   ARITHMETIC. 

CHAPTER   I. 

INTRODUCTION. 

1.  Matter  is  anything  that  occupies  space. 

2.  Divisions   of   Matter.     Three    divisions   of  matter 
are  recognized  in  science,  —  masses,  molecules,  and  atoms. 

A  mass  of  matter  is  any  portion  of  matter  appreciable 
by  the  senses. 

A  molecule  is  the  smallest  particle  of  matter  into  which 
a  body  can  be  divided  without  losing  its  identity ;  or  it  is 
the  smallest  portion  of  matter  which  can  exist  by  itself. 

An  atom  is  a  still  smaller  particle  produced  by  the 
division  of  a  molecule;  or  it  is  the  smallest  portion  of 
matter  that  can  go  into  combination. 

EXAMPLES.  The  sun  and  a  grain  of  sand  are  masses  of 
matter.  The  smallest  particle  of  salt  which  can  exist 
and  which  exhibits  the  properties  of  salt  is  a  molecule. 
The  minute  particles  of  chlorine  and  sodium  which  com- 
pose the  molecule  of  salt  are  atoms. 

A  mass  is  made  up  of  molecules,  and  a  molecule  is  com- 
posed of  atoms. 

3.  Attractions  of  Matter.     The  three  forms  of  attrac- 
tion admitted  in  science  are  :  — 


2  CHEMICAL   ARITHMETIC. 

First.    Gravitation,  or  the  attraction  between  masses. 

Second.  Cohesion,  or  the  attraction  between  like  mole- 
cules; adhesion  between  unlike  molecules. 

Third.  Chemical  attraction,  or  the  attraction  between 
unlike  atoms. 

EXAMPLES.  The  attraction  between  the  sun  and  the 
planets,  or  between  the  earth  and  all  bodies  upon  it,  is 
gravitation.  The  attraction  between  the  molecules  of  a 
piece  of  marble  is  cohesion.  The  attraction  between  a 
liquid  and  solid,  —  as,  for  instance,  when  you  dip  your 
hand  into  water  it  becomes  wet,  —  or  between  two  different 
solids  at  the  surface,  as  shown  by  the  action  of  cements, 
is  adhesion.  The  attraction  between  the  unlike  atoms  of 
chlorine  and  sodium  by  means  of  which  we  have  an 
entirely  different  substance,  salt,  is  chemical  attraction. 

4.  Province  of  Physics.     Physics  is  that  department 
of  physical  science  which  studies  the  results  which  come 
from  the  molar  and  molecular  conditions  of  matter. 

5.  Province  of  Chemistry.     Chemistry  studies  matter 
in  its  atomic  condition.     It  investigates  the  laws  and  con- 
ditions  of   chemical  changes,   and   seeks   to    account  for 
some  of  the  phenomena  connected  therewith. 

6.  Physical    Changes.       Physical    changes    are    those 
which   take   place    outside   the  molecule;    they  have   no 
effect  upon  the  molecule  itself  nor  alter  the  identity  of 
the  matter  operated  on.    The  study  of  physics  is  a  study  of 
physical  changes. 

7.  Chemical  Changes.     Chemical  changes  take  place 
through  the  atoms  and  within  the  molecule.     They  alter 
the  character   of  the    molecule,   and   hence   destroy   the 


INTRODUCTION.  3 

identity  of  the  matter  itself.     The  study  of  chemistry  is  a 
study  of  chemical  changes. 

EXAMPLES.  The  change  of  water  into  ice  and  steam, 
or  the  change  of  any  solid  into  a  liquid,  or  of  any  liquid 
into  a  vapor,  are  physical  changes.  But  when  water  is 
subjected  to  the  influence  of  the  electric  current,  it 
undergoes  a  more  radical  change ;  the  water  disappears, 
and  in  its  place  appear  two  gaseous  substances,  oxygen  and 
hydrogen,  entirely  different  from  the  water  from  which 
they  were  derived.  This  is  a  chemical  change. 

8.  Physical  Properties.     Physical  properties  are  those 
properties  which  bodies  possess  in  virtue  of  their  molar 
or  molecular  condition. 

9.  Chemical    Properties.       Chemical    properties     are 
those  which  result  from  the  atomic  composition  of  the 
molecule. 

1C.  Chemistry  defined.  Chemistry  is  that  branch  of 
physical  science  which  treats  of  the  atomic  composition  of 
bodies,  and  of  those  changes  in  matter  which  result  from 
an  alteration  in  the  kind,  the  number,  or  the  relative 
position  of  the  atoms  which  compose  the  molecule. 

11.  Analysis  and  Synthesis.  The  two  processes  by 
which  the  chemist  seeks  to  find  out  the  composition  of 
matter  are  analysis  and  synthesis. 

Analysis  consists  in  separating  the  molecule  into  its 
constituent  atoms. 

Synthesis  consists  in  putting  together  constituent  atoms 
to  form  the  molecule. 


CHEMICAL   ARITHMETIC. 


CHAPTER   II. 

MOLECULES   AND   ATOMS. 

12.  Chemical  Definition  of   the  Molecule.     A  mole- 
cule is  the  smallest  particle  of  any  substance  which  can 
exist  in  a  free  state  in  nature. 

Molecules  classified.     Molecules  are  of  two  classes :  — 
First.  Elementary  molecules,  or  those  whose  atoms  are 

alike. 

Second.   Compound  molecules,  or  those  whose  atoms  are 

unlike. 

13.  Simple  Substances  are  those  whose  molecules  con- 
tain like  atoms. 

14.  Compound  Substances  are  those  whose  molecules 
contain  unlike  atoms. 

15.  Number  of  Simple  Substances.     There  are  sixty- 
eight  elementary  substances,  as  far  as  has  been  investi- 
gated by  chemical  science  |  that  is,  sixty-eight  substances 
whose   molecules   contain    like   atoms.      Therefore   it  is 
obvious  that  there  are  sixty-eight  different  kinds  of  atoms. 
From  combinations  of  these  sixty-eight  kinds  of  atoms  all 
the  different  varieties  of  matter  result.     We  cannot  resolve 
a  simple  substance  into  any  other  substances  or  atoms. 

16.  Ampere's    Law.      "Equal   volumes   of   all  gases, 
simple  as  well  as  compound,  under  like  conditions  of  tem- 
perature and  pressure,  contain  the  same  number  of  mole- 
cules." 


MOLECULES    AND   ATOMS.  <"  5 

From  this  law,  which  is  the  most  -important  law  of 
modern  chemistry,  it  results,  — 

First.  That  the  molecules  of  all  bodies  in  the  gaseous 
state  are  of  the  same  size. 

Second.  That  the  weight  of  any  molecule,  compared 
with  that  of  hydrogen,  is  proportional  to  the  weight  of 
any  given  volume,  also  compared  with  the  same  volume  of 
hydrogen. 

EXAMPLES.  If  1  liter  of  nitrogen,  which  weighs  14 
times  as  much  as  a  liter  of  hydrogen,  contains  the  same 
number  of  molecules,  then  it  is  obvious  that  each  molecule 
of  nitrogen  must  be  14  times  as  heavy  as  a  molecule  of 
hydrogen. 

1 7.  Number  of  Atoms  in  the  Molecule  of  Hydrogen. 

Assuming  that  1  volume  of  hydrogen  contains  1000 
molecules,  then,  according  to  the  law  of  Ampere,  1  vol- 
ume of  chlorine  will  contain  1000  molecules  also. 

Suppose  these  volumes  (that  is,  1  volume  of  hydrogen 
containing  1000  molecules  and  1  volume  of  chlorine  con- 
taining 1000  molecules)  be  mixed  together  and  exposed 
to  the  action  of  the  sunlight,  they  combine,  forming  2 
volumes  of  the  new  substance,  hydrochloric  acid  gas, 
which  2  volumes,  by  the  same  law,  will  contain  2000 
molecules.  Upon  analysis,  each  molecule  of  hydrochloric 
acid  gas  will  be  found  to  contain  1  atom  of  hydrogen 
and  1  atom  of  chlorine.  That  is,  the  2000  molecules 
will  contain  2000  atoms  of  hydrogen  and  2000  atoms  of 
chlorine.  The  2000  molecules  will  contain,  therefore, 
4000  atoms ;  or,  each  molecule  will  contain  2  atoms. 
Hence  each  molecule  of  hydrogen  is  made  up  of  2 
atoms. 


6  CHEMICAL   ARITHMETIC. 

18.  Molecular  Weights/     If  the  weight  of  the  hydro- 
gen atom  be  taken  as  1,  then,  since  its  molecule  contains 
2  atoms,  its  molecular  weight  will  be  2. 

Since  the  molecule  of  a  compound  gas  or  vapor  occupies 
a  volume  twice  as  large  as  that  occupied  by  the  atom  of 
hydrogen,  it  is  obvious  that  the  specific  gravity  of  the  gas 
or  vapor  may  be  found  from  the  molecular  weight  by 
dividing  the  latter  by  2.  The  specific  gravity  of  a  com- 
pound gas  or  .vapor  is,  therefore,  one-half  its  molecular 
weight.  The  molecular  weight  of  any  substance  may  be 
obtained  by  multiplying  its  density  in  the  state  of  gas  by 
the  molecular  weight  of  hydrogen  ;  that  is,  by  2. 

EXAMPLES.  The  density  of  oxygen  gas,  for  example,  is 
16 ;  that  is,  any  volume  as  1  liter  weighs  16  times  as 
much  as  1  liter  of  hydrogen.  Its  molecule  must  be, 
therefore,  16  times  as  heavy.  The  molecular  weight  of 
hydrogen  is  2 ;  therefore  the  molecular  weight  of  oxygen 
will  be  16  x  2  =  32. 

The  weight  of  1  liter  of  hydrogen  is  called  1  crith, 
and  the  weight  of  the  hydrogen  atom  1  microcrith. 

19.  Number  of  Atoms  in  the  Molecule.     The  number 
of  atoms  in  a  molecule  is  obtained  by  dividing  the  molecu- 
lar weight  by  the  atomic  weight. 

EXAMPLE.  The  molecular  weight  of  oxygen  is  32,  and 
its  atomic  weight  16.  The  number  of  atoms  in  the  mole- 
cule is  32  divided  by  16  =  2.  The  molecular  weight  of 
phosphorus  is  124,  and  its  atomic  weight  31;  its  mole- 
cule, therefore,  contains  4  atoms. 


MOLECULES  AND  ATOMS.  7 

PROPERTIES  OF  ATOMS. 

20.  Definition.      An  atom  is  the  smallest  particle  of 
simple  matter  which  can  enter  into  the  composition  of  a 
molecule. 

21.  Atomic  Weight.    The  relative  weight  of  any  atom 
referred  to  hydrogen  as  unity  is  its  atomic  weight.     It  is 
the  smallest  weight  of  any  simple  substance  which  can 
take  part  in  the  formation  of  a  chemical  compound. 

The  molecular  weight  of  any  substance  is  the  sum  of 
the  weights  of  its  constituent  atoms. 

22.  Quaiitivalence.     The  quantivalence  of  an  atom  is 
the  quality  of  its  combining  power,  expressed  in  hydrogen 
units.     It  expresses  the  number  of  hydrogen  atoms  with 
which  it  can  combine  or  for  which  it  can  be  exchanged. 

EXAMPLES.  The  quantivalence  of  zinc  is  2,  because 
1  atom  replaces  2  of  hydrogen.  The  quantivalence 
of  carbon  is  4,  because  1  atom  of  carbon  requires 
4  of  hydrogen  to  satisfy  it  in  combination.  Atoms  are 
called  monads,  dyads,  triads,  tetrads,  pentads,  hexads,  and 
heptads,  according  to  their  quantivalence.  The  Latin  nu- 
merals are  used  for  the  adjective  terms.  These  atoms  are 
univalent,  bivalent,  trivalent,  quadrivalent,  quinquivalent, 
sexivalent,  and  septivalent.  Atoms  whose  quantivalence 
is  even  are  called  artiads ;  those  whose  quantivalence  is 
odd  are  called  perissads. 

An  atom  may  form  several  compounds  with  the  same 
substance.  Therefore  its  quantivalence  may  vary.  It 
always  increases  or  diminishes  by  2,  so  that  it  may  have 
quantivalence  of  1,  3,  5,  or  7,  or  of  2,  4,  or  6.  A  perissad 
atom  can  never  become  an  artiad  by  such  a  change,  nor 
can  an  artiad  become  a  perissad. 


8  CHEMICAL   ARITHMETIC. 

EXAMPLES.  'Iron  in  iron  sulphate  is  a  dyad,  in  pyrites 
it  is  a  tetrad,  and  in  ferric  acid  a  hexad.  Chlorine  forms  a 
series  of  compounds  with  oxygen  in  which  its  quantivalence 
is  1,  8,  5,  and  7. 

Atoms  are  divided  into  two  classes,  according  to  the 
quality  of  their  combining  power. 

First.  Positive  atoms  are  those  which  are  attracted  to  the 
negative  pole  in  electrolysis,  and  whose  hydrates  are  bases. 

Second.  Negative  atoms  are  those  which  go  to  the  posi- 
tive electrode,  and  whose  hydrates  are  acids. 

23.  Atomic  Symbols.     Berzelius,  in  1815,  proposed  an 
abbreviated  form  of  chemical  language.      In  this  system 
each  atom  has  for  its  symbol  the  first  letter  of  its  Latin 
name.     When  the  names  of  two  different  atoms  begin  with 
the  same  letter,  a  second  letter  suggestive  of  the  name  is 
added. 

EXAMPLES.  Ag  stands  for  an  atom  of  silver ;  Fe,  for  an 
atom  of  iron ;  Sn,  for  one  of  tin,  etc.  (on  page  89  will  be 
found  the  table  of  the  symbols  of  the  elements). 

Each  atomic  symbol  stands  not  only  for  the  atom,  but 
represents  its  atomic  weight. 

EXAMPLES.  Fe  (ferrum)  represents  56  weight-units  of 
iron;  Hg  (hydrargyrum),  200  weight-units  of  mercury; 
O,  16  weight-units  of  oxygen. 

24.  The  quantivalence  of  an  atom  is  indicated  by  plac- 
ing Roman  numerals  above  or  a  little  to  the  right  of  the 
symbol.     Sometimes  minute-marks  are  used. 

EXAMPLES.  1.  H  or  H'  stands  for  the  monad  hydrogen 
atom;  2.  S  or  S"  stands  for  the  bivalent  sulphur  atom; 
3.  P  or  P";  for  the  trivalent  phosphorus  atom;  4.  C  or 
C'r"  for  the  quadrivalent  carbon  atom. 


MOLECULES   AND   ATOMS.  9 

Sometimes  graphic  symbols  are  used  to  represent  the 
quantivalence  atoms,  the  graphic  symbols  being  a  circle 
with  lines  called  bonds  radiating  from  it ;  as,  for  example, 

Monad.         Dyad.  Triad.  Tetrad.        ^  Pentad.          Hexad. 

6     -o-     A     -9-    ')6;     £ 

The  circles  are  usually  omitted,  the  bonds  radiating  from 
the  symbol.  The  number  of  bonds  and  not  their  direc- 
tion is  significant,  as,  for  example,  -Q-  Q-  O~  stands 

I 
equally    for    1    atom    of    dyad   oxygen.      N=,    N=,   or 

— N=  equally  represent  the  atom  of  trivalent  nitrogen. 

25.  Multiplication  of  Atoms.  Atoms  are  multiplied 
by  placing  an  Arabic  numeral  below  and  to  the  right  of 
the  symbol. 

EXAMPLES.  C2  represents  2  atoms  of  carbon.  N*,  4 
atoms  of  nitrogen.  C13,  3  atoms  of  chlorine. 

Molecules  are  multiplied  by  enclosing  their  symbols  in 
brackets  and  placing  the  numeral  outside,  below,  and  to 
the  right. 

EXAMPLES.  (H2)6  represents  6  molecules  of  free 
hydrogen.  (Br2)2  stands  for  2  molecules  of  bromine. 


10  CHEMICAL   ARITHMETIC. 


CHAPTER  III. 

COMPOUND   MOLECULES   AND   VOLUME  RELATIONS. 

26.  Compound   Molecule.     A   compound   molecule  is 
one  whose  constituent  atoms  are  unlike.     Compound  mole- 
cules are  formed  by  the  union  of  atoms  according  to  the 
law  of  quantivalence. 

27.  Molecular  Weight.      The    molecular   weight  of  a 
compound  molecule  is  the  sum  of  the  atomic  weights  of  its 
constituents.     It  is  always  equal  to  twice  the  density  of 
the  substance  in  the  state  of  gas. 

28.  Classification   of   Compound    Molecules.      Com- 
pound molecules  are  divided  into  two  classes :  first,  those 
whose  atoms  are  directly  united,  called  Binaries ;  second, 
those  whose  atoms  are  indirectly  united,  called  Ternaries. 
A  binary  compound  is  formed  by  the  union  of  two  simple 
substances,  the  termination  IDE  being  the  characteristic : 
as,  for  example,  sodium  and  chlorine  yield  sodium  chloride; 
silver  and  sulphur  yield  silver  sulphide  ;  calcium  and  iodine 
yield  calcium  iodide.     In  some  cases  the  number  of  atoms 
of  each  constituent  is  to  be  indicated. 

This  is  done  by  prefixing  Greek  numerals  to  each  of  the 
names  given ;  as,  for  example,  1  atom  of  C  and  2  of  O  form 
carbon  dioxide,  1  atom  of  P  and  5  of  Br  form  phosphorus 
pentebromide. 

29.  Definition  of  an  Acid.     An  acid  molecule  is  one 
which  consists  of  one  or  more  negative  atoms  united  by 


COMPOUND   MOLECULES   AND   VOLUME  RELATIONS.     11 

hydrogen  and  Coxygsib  It  is  a  compound  of  hydrogen 
and  oxygen  with  some  non-metallic  element,  and  possesses 
the  property  of  turning  blue  litmus  paper  or  solution  red. 

30.  Definition  of  a  Base.      A  basic  molecule  is  one 
which    contains    one  or  more  positive  atoms   united   by 
hydrogen  and  oxygen.     It  is  a  compound  of  hydrogen, 
oxygen,  and   some    metallic   element,  and   possesses   the 
property  of  restoring  the  color  to  vegetable  blues,  which 
have  been  reddened  by  an  acid. 

31.  Definition  of  a  Salt.     A   saline    molecule   is   one 
which  contains  a  positive  atom  or  group  of  atoms,  united 
by  oxygen  to  a  negative  atom  or  group  of  atoms.     It  is 
formed  by  the  action  of  an  acid  upon  a  base,  and  since  it 
contains  no  hydrogen,  has  no  action  upon  vegetable  colors. 

32.  Compound  Radical.      A    compound  radical   is   a 
group  of  atoms,  which  goes  into  combination  like  a  single 
atom.     It  may  be  composed  of  two  or  more  elements;  as, 
for  example,  (NH4)  ammonium,  (C2H5)  ethyl. 

33.  Normal,  Acid,  Basic,  and  Double  Salts.     A  salt 
is  formed  by  substituting  a  metal  for  the  hydrogen  of  an 
acid,  each  bond  of  the  metal  displacing  one  atom  of  hydro- 
gen.    A  normal  salt  is  formed  by  displacing  all  the  hydro- 
gen of  the  acid  with  an  equivalent  metal.     An  acid  salt  is 
formed  by  exchanging  a  part  of  the  hydrogen  of  an  acid  for 
an  equivalent  of  metal. 

A  basic  salt  is  formed  by  the  substitution  of  a  metal  in 
part  for  the  hydrogen  of  an  acid,  and  in  part  for  the  half 
or  the  whole  of  the  hydrogen  of  water  (H2O). 

Double  salts  are  those  containing  two  or  more  different 
positive  or  metal  atoms. 


12  CHEMICAL   ARITHMETIC. 

EXAMPLES.  K'  with  HNO3  forms  KNO3,  displacing  H. 
KNO3  is  a  normal  salt,  formed  also  by  acting  upon  HNO3 
by  KHO,  as  :  - 

HNO3  -f  KHO  =  KNO3  +  H2O. 

K'  with  H2SO4  may  form  KHSO4,  an  acid  sulphate  formed 
also  by 

KHO  +  H2SO4  =  HKSO4  +  H2O. 


Bi'"  with     ga  forms  Bi  |  Q°3,  usually  written  BiONO3, 

a  basic  nitrate. 

NaCa"SbO4,  sodio-calcium  antimonate  is  an  example  of 
a  double  salt,  or  Ba"Zn"SiO4,  baro-zincic  silicate. 

Monobasic  acids  can  form  only  normal  salts.  Polybasic 
acids  can  form  normal,  acid,  and  double  salts. 

34.  Chemical  Equations.     A  chemical  equation  is  the 
expression  in  symbols  of  a  chemical  reaction,  or  change. 
The  sign  plus  (+)  indicates  added  to,  and  the  sign  minus 
(—  ),  taken  from,  and  the  sign  of  equality  (  =  ),  equals  to. 
The  equation  must  be  a  true  equation  ;  that  is,  the  sum  of 
the  weights  of  the  atoms  on  one  side  must  equal  the  sum 
of  the  weights  of  the  atoms  on  the  other  side. 

35.  The   substances    entering    into    the    reaction    are 
called  factors;  these  constitute  the  first  member.      The 
substances  issuing  from  the  reaction  are  called  products; 
these  constitute  the  second  members. 

The  equation,  representing  the  reaction  of  two  molecules 
upon  each  other  may  be  written  by  the  following  rule  :  - 

Place  the  formulas  of  the  factors,  connected  by  the 
sign  plus,  as  the  first  member  of  the  equation,  and  the 
formulas  of  the  products,  also  connected  by  the  sign 
plus,  as  the  second. 


COMPOUND   MOLECULES    AND    VOLUME   RELATIONS.     13 

EXAMPLES.  Let  AB  and  EF  be  two  molecules.  The 
reaction  between  them  would  be  represented  by  the 

equation 

AB  +  EF=  AF  +  BE. 

36.  Weight   of    the    Factors    and    Products.       The 

quantities  of  matter  taking  part  in  a  chemical  change  are 
definite  in  weight,  since  each  formula  represents  a  definite 
weight,  viz.,  the  molecular  weight.  For  the  same  reason 
no  loss  of  weight  can  be  the  result  of  any  chemical  reac- 
tion. 

37.  There  are  three  kinds  of  chemical  reactions  :  — 
First.  Analytical  reactions  ;  that  is,  the  separation  of  a 

complex  molecule  into  simpler  ones. 

Second.  Synthetical  reactions,  or  the  union  of  two  or 
more  simple  molecules  to  form  a  more  complex  one. 

Third.  Metathetical  reactions,  or  the  transposition  or 
exchange  of  atoms  between  molecules. 

EXAMPLES.  An  analytical  reaction  may  be  represented 
by  the  general  equation 


or,  taking  an  actual  example, 

HgO    =    Hg    +    O; 

Mercuric  oxide.    Mercury.    Oxygen. 

that  is,  one  molecule  of  mercuric  oxide  will   yield  one 
molecule  of  mercury  and  one  molecule  of  oxygen. 

Synthetical  reactions  may  be  represented  by  the  general 

equation 

E+F=EF\ 

or,  taking  an  actual  example, 

Fe  +  S    =    FeS; 

Iron.    Sulphur.  Iron  sulphide. 


14  CHEMICAL   ARITHMETIC. 

that  is,  one  molecule  of  iron  and  one  molecule  of  sulphur 
yield  one  molecule  of  iron  sulphide. 

Metathetical  reactions  may  be  represented  by  the  gen- 
eral formula 

AB  -f  EF=  AF+  BE ; 
or,  practically, 

Zn  +  H2SO4    =     ZnSO4    +    2H. 

Zinc  +  Sulphuric  acid  =  Zinc  sulphate  -*-  TT^     -ogen. 

38.  The  conditions  which  form  chemical  change  depend 
upon  the  facility  with  which  the  atoms  of  any  molecule 
may  be  rearranged.  When  substances  are  in  the  gaseous 
or  liquid  state,  these  changes  between  atoms  take  place 
most  readily.  Hence,  fusion  or  solution  or  vaporization 
facilitate  chemical  action.  Heat  is  therefore  the  great  aid 
to  the  chemist. 


ST01CHIOMETRY.  15 


CHAPTER  IV. 

(  STOICHIQMETBY.  ]  —  CHEMICAL   ARITHMETIC. 

39.  Definition."    By  Stoichiometry  we   mean  that  de- 
partment of  chemistry  which  treats  of  the  numerical  rela- 
tions  of    atoms.      The    calculations   of    these   numerical 
relations,  whether  of  volume  or  weight,  depend  upon  the 
fact  that  every  atom  has  its  own  weight,  called  the  atomic 
weight.     The  atomic  weight  is  the  smallest  portion  by 
weight  of  any  simple  or  elementary  substance  —  referred 
to  the  atom  of  hydrogen  as  unity  —  which  can  take  part 
in  a  chemical  change. 

40.  Unit   of    Weight.     The   weight   of   the  hydrogen 
atom  is  called  a  microcrith.     (The  weight  of  one  liter  of 
hydrogen  under  general  conditions  of   temperature  and 
pressure  is  one  crith.)     We  adopt  the  term  microcrith  for 
convenience'  sake. 

41.  All  chemical  changes  take  place  between  definite 
quantities  of  matter,  as  represented  by  a  chemical  equa- 
tion.    An  equation  expresses  not  only  the  fact  of  chemical 
reaction  between  two  bodies,  but  also  indicates  the  quan- 
tities by  weight  concerned  in  it. 

RULES. 

42.  From  the  Formula  of  a  Substance   to   find   its 
Molecular  Weight.      The  molecular  weight   of  a   com- 
pound is  the  sum  of  the  atomic  weights  of  all  the  atoms  of 
the  elements  which  compose  it. 


16  CHEMICAL   AEITHMETIC. 

i 

The  name  of  each  element  present  being  written  in  a 
column,  and  opposite  to  each  the  multiple  of  its  atomic 
weight  which  is  present  in  the  compound,  on  adding  these 
numbers  together  the  molecular  weight  of  the  compound 
is  obtained. 

Thus  the  molecular  weight  of  sulphuric  acid,  H2SO4,  is 

H=    1x2=2 

S~=32x  1  =  32 

O  =  16x4  =  64 

98 

43.  To  find  the  Percentage  Composition  of  any  Sub- 
stance in  a  Molecule.  Rule.  Multiply  the  atomic  weight 
by  the  number  of  atoms,  and  this  product  by  100.  Divide 
the  final  product  by  the  molecular  weight,  and  the  quo- 
tient will  be  the  percentage  amount  of  that  constituent. 

EXAMPLE.  What  is  the  percentage  composition  of 
carbon  dioxide,  Co2? 

Carbon  =12 

Oxygen,  16  x  2  =  3? 
Molecular  weight,  44 

Carbon  =  12  x  -^  =  27.27  per  cent. 
Oxygen  =  32  x  -^  =  72.73  per  cent. 

This  rule  can  be  expressed  by  a  general  formula.  Let 
m  represent  the  molecular  weight,  a  the  atomic  weight  of 
any  constituent,  n  the  number  of  atoms,  and  x  its  percen- 
tage amount  ;  then  we  have  the  proportion  :  — 

tnfian:  :  100  :#, 
from  whence  the  formula 


m 


STOICHIOMET&Y.  17 

In  the  above  formula,  when  any  three  of  the  quanti- 
ties a,  w,  w,  and  x  are  known,  the  fourth  can  be  found. 
Whence,  to  find  the  number  of  atoms  of  any  constituent  in 
a  molecule,  "a?,"  "a,"  and  "w"  being  known,  we  have,  by 
transposing  formula  (1)  :  - 

n  =  -^L,  (2) 

alS°  a=   m*  (3) 

am*  m  =  an  X  1QQ  (4) 

44.    To  calculate  from   an  Equation  a  Mass. 

Mule.  Find  the  multiples  of  the  atomic  or  molecular 
weights  of  the  substances  given  and  asked  in  the  equation, 
and  work  the  proportion. 

The  molecular  weight  of  substance  given  :  the  molecu- 
lar weight  of  substance  asked  :  :  the  real  mass  of  sub- 
stance given  :  the  real  mass  of  substance  asked.  Thus, 
to  find  now  many  grams  of  sodium  sulphate  can  be  ob- 
tained from  100  grams  of  sodium  hydrate  :  — 

EXAMPLE  1.     2  NaHO  +  H2SO4  =  2  H2O  +  Na2SO4. 
2  x  40  142. 

2  X  40  :  142  : :  100  :  x 
x  =  J-4JMLO  =  177.5  grams. 

EXAMPLE  2.     In  the  equation 

KNO3  +  H2SO4  =  HNO3  +  HKSO4 
101    +     98     =     63     +    136. 

125  grams  of  KNO3  yield  77.97  grams  of  HNO3,  whose 
molecular  weight  is  63.  What  is  the  molecular  weight 
of  KN(X? 


18  CHEMICAL  ARITHMETIC. 

This  rule  can  be  simply  expressed  by  the  general  pro- 
portion :  — 

M  :  tn  :  :  W:  w, 

where  M  represents  the  molecular  weight  of  the  substance 
given,  m  the  molecular  weight  of  the  substance  asked,  W 
the  real  mass  of  the  substance  given,  and  w  the  real  mass 
of  the  substance  asked  ;  whence 


45.  The  Relations  of  Weight  to  Volume.     1.    To  find 
the  volume  occupied  by  a  given  weight  of  any  gas.     Rule. 
Divide  the  weight  of  the  gas  given  by  the  weight  of  1 
liter  ;  the  quotient  is  the  number  of  liters. 

2.  To  find  the  weight  of  any  given  volume  of  gas. 
Rule.  Multiply  the  number  of  liters  of  gas  by  the  weight 
of  1  liter;  the  product  is.  the  weight  of  the  given  volume. 

EXAMPLES.  1.  What  volume  is  occupied  by  6.08  grams 
of  oxygen,  the  weight  of  1  liter  of  oxygen  being  1.43 
grams  ?  6.08  -=-  1.43  =  4.25  liters.  Am. 

2.  What  is  the  weight  of  25  liters  of  nitrogen  gas,  1 
liter  weighing  1.26  grams?  1.26  x  25  =  31.5  grams.  Ana. 

46.  Density  of  Gases.      The  density  of  any  gas  ex- 

presses how  many  times  the  gas  is  heavier  than  hydrogen. 
Knowing  the  density,  the  weight  of  1  liter  may  readily 
be  obtained  by  multiplying  it  by  the  weight  of  1  liter  of 


STO1CHIOMETBY.  19 

hydrogen,  0.0896  grams,  or  1  crith.  The  molecular  weight 
of  any  substance  being  the  weight  of  2  volumes  of  the 
substance  in  the  state  of  gas,  it  is  evident  that  its  density 
in  the  state  of  gas  may  be  obtained  by  dividing  its  molec- 
ular weight  by  2.  With  few  exceptions,  the  density  of 
any  elementary  gas  is  expressed  by  the  same  number  as  its 
atomic  weight,  and  that  of  any  compound  gas  is  expressed 
by  the  same  number  as  half  its  molecular  weight.  Thus, 
oxygen,  O  =  16  ;  density,  16  ;  or  1  liter  weighs  16  criths. 
Ammonia,  NH3  =  17  ;  density,  8.5  ;  or  1  liter  weighs  8.5 
criths. 

47.  Relation  of  Gaseous  Volume  to  Pressure.  To 
calculate  the  change  in  volume  of  a  mass  of  gas  produced 
by  a  change  in  pressure. 

Boyle's  Law.  The  volume  of  a  mass  of  gas  varies  in- 
versely as  the  pressure  upon  it;  or  the  volume  of  a  mass 
of  gas,  multiplied  by  the  pressure  at  any  one  time,  is  equal 
to  the  volume  of  the  same  mass  of  gas  multiplied  by  the 
pressure  upon  it  at  any  other  time.  Thus,  let  V  equal  the 
volume  of  a  gas  under  the  pressure  P,  and  let  V  equal  the 
volume  under  the  pressure  P'  ;  then 

' 


,  or  F= 

If  the  pressure  upon  1000  cc.  of  gas  be  increased  from 
400  mm.  to  800  mm.,  what  is  the  new  volume  ? 


48.    Relation    of    Gaseous  Volume  to    Temperature. 

Gruy  Lussac's  Law.  When  273  volumes  of  gas  at  0°  C.  are 
heated,  they  increase  by  one  volume  for  every  1°  C.  through 
which  thev  are  heated.  Thus  :  — 


20  CHEMICAL  ARITHMETIC. 

273  volumes  of  gas  at  0°  C.  become  at  1°  C.  273  +  1  volumes, 

273        "     .         "          "  "        "  2°  C.  273  +  2         " 

273        "  "          "  "         "  3°  C.  273  +  3         " 

273        "  "  "  "         u  t°  C.  273  +£          " 

where  t  expresses  any  number  of  degrees  on  the  centigrade 
scale. 

The  coefficient  of  the  expansion  of  a  gas  is  ^  of  the 
volume  of  the  gas  at  0°  for  every  degree  centigrade.  Hence 
v  volumes  at  t°  C.  become  at  T°  C. 

..      273  +  T 


volumes  ;  which,  if  V  stands  for  the  volume  of  the  gas  after 
change  of  temperature  t°  C.  to  T°  C.,  is  usually  written:  — 


273  +* 

EXAMPLE  1.  Find  the  new  volume,  if  1000  cc.  of  gas 
are  heated  from  17°  C.  to  27°  C.  The  formula  is  :  - 

1000  (273  +  27)  =  1000  X  300  =  1(m  g  c 
273  +  17  290 

EXAMPLE  2.  If  1000  cc.  of  gas  at  -  23°  C.  are  heated 
to  27°  C.,  find  the  new  volume. 

F=  1000  (273  +  27)  =  1000  X  300  =  12Q()  cc 
273  -  23  250 

49.  If  the  pressure  on  the  gas,  as  well  as  its  tempera- 
ture, be  changed,  the  above  formula  must  be  combined 
with  the  one  given  in  (46). 


273  +  *          P' 


J 


STOICHIOMETBY.  21 


EXAMPLE.  If  -500  cc.  of  gas  are  cooled  from  39°  C.  to 
13°  C.,  the  pressure  being  decreased  from  800  mm.  to 
300  mm.,  find  the  new  volume. 

500(273  +  13)      800  =  1222<2cc> 
273  +  39  300 

5O.  Density  of  Oases.  When  the  temperature  of  and 
the  pressure  on  a  gas  are  not  mentioned,  it  is  supposed 
to  be  at  760  mm.  and  0°  C.  A  gas  under  these  conditions 
is  said  to  be  normal. 

The  formulae  given  in  this  and  the  succeeding  section 
only  apply  to  normal  gases  ;  hence,  when  necessary,  the 
gas  under  consideration  must  be  rendered  normal  by  using 
the  formula  :  — 

F==  0X273      P' 
273+*       760' 

and,  conversely,  the  volume  found  by  these  formulas  is  nor- 
mal, and  must  be  reduced  to  the  required  temperature  and 

pressure  by 

T        760 


273  m 

In  the  case  of  gases,  the  liter,  =  1000  cc.,  is  taken  as  the 
unit  of  volume,  and  the  mass  of  one  liter  of  normal  hydro- 
gen, called  a  crith,  =  .0896  gram,  is  taken  as  the  unit  of 
mass. 

The  density  of  a  gas,  then,  is  the  number  of  criths  con- 
tained in  one  liter  of  it,  measured  at  0°  C.  and  760  mm.  ; 
or  the  number  of  times  it  is  heavier  than  an  equal  volume 
of  hydrogen. 

Hence  the  mass  in  grams  of  a  liter  of  any  normal 
gas  can  be  found  by  multiplying  its  density  by 
.0896. 


22  CHEMICAL   ARITHMETIC. 

EXAMPLE.  The  density  of  carbon  monoxide  is  14 ;  re- 
quired the  weight  of  one  liter. 

14  x  .0896  =  1.2544  grams. 

The  density  of  a  gas  referred  to  air  may  be  obtained  by 
multiplying  its  density  referred  to  hydrogen  by  .06926,  the 
density  of  hydrogen  referred  to  air. 

EXAMPLE.  Nitric  oxide  is  15  times  as  heavy  as  hydro- 
gen ;  how  many  times  is  it  heavier  than  air  ? 

15  X  .06926  =  1.0389. 

If  the  density  of  a  gas  referred  to  air  be  given,  its  den- 
sity referred  to  hydrogen  can  be  obtained  by  multiplying 
its  density  referred  to  air  by  14.436. 

EXAMPLE.  Sulphur  dioxide  is  2.22  times  as  heavy  as 
air ;  find  its  density  and  molecular  weight. 

14.436  x  2.22  =  32.042 
is  the  density  referred  to  hydrogen,  and 

32.042  x  2  =  64.084 
is  the  molecular  weight. 

51.  Volume  and  Mass  of  Gases.  It  is  found  by  ex- 
periment that  22.32  liters  of  any  normal  gas  weigh  a  num- 
ber of  grams  equal  to  the  number  expressing  the  molecular 
weight  of  the  gas.  Thus  :  — 

22.32  liters  of  hydrogen  (H2  =    2)  weigh    2  grams. 
22.32      "     "  oxygen       (O2  =  32)      "      32      " 
22.32      "     "  nitrogen    (N2  =  28)      "      28      " 
22.32      "     "chlorine    (C12  =  71)      "      71      "' 

This  volume,  22.32  liters,  is  commonly  spoken  of  as  "two 
volumes  "  and  expressed  by  the  symbol  CD. 


STOIC  HIOMETBY.       f^y  23 

Since  22.32  liters  (or,  if  great  accuracy  be  not  required, 
22.4  liters)  of  any  gas  weigh  its  molecular  weight  in 
grams,  a  liter  of  any  gas  weighs  its  molecular  weight  in 
grams  divided  by  22.32  (or  22.4);  and  one  gram  of  any 
gas  occupies  22.32  (or  22.4)  liters  divided  by  its  molecular 
weight.  Hence  the  mass  in  grams  of  any  volume  of  a  gas 
can  be  found  by  multiplying  the  number  of  liters  of  it  by 
its  molecular  weight,  and  dividing  by  22.32.  - 

EXAMPLE  1.  Find  the  mass  of  250  liters  of  chlorine 
(C12  =  71). 


Conversely,  the  volume  in  liters  of  any  gas  can  be  found 
by  multiplying  the  number  of  grams  of  it  by  22.32,  and 
dividing  by  the  molecular  weight. 

EXAMPLE  2.  Find  the  volume  of  225  grams  of  hydro- 
gen sulphide  (H2S  =  34). 

225  X  22.32  =  147^rs> 
34 

EXAMPLE  3.  Find  the  mass  of  80  liters  of  oxygen 
(O2  =  32)  measured  at  52°  C.  and  7.40  mm.  The  gas  must 
be  reduced  to  0°  C.  and  760  mm. 


80X273  X  ™  :  x    **-  =  =  93.47 

273  +  52      760      22.4        95 


52.  Equation  and  Volumes  of  Gases.  When  the 
volume  of  one  gas  is  given,  and  that  of  another  gas  is 
asked,  since  each  molecular  weight  expresses  two  volumes 
of  the  gas,  the  result  may  often  be  obtained  directly. 


24  CHEMICAL   ARITHMETIC. 

EXAMPLE  1.  What  volume  of  hydrogen  chloride  is 
formed  when  10  liters  of  chlorine  combine  with  hydrogen? 

C12  +  H2=2HC1. 

j-ri+rTj  =  2  m 

Two  volumes  of  chlorine  form  twice  two  volumes  of 
hydrogen  chloride  ;  hence  10  liters  of  chlorine  form  2  x  10 
=  20  liters  of  hydrogen  chloride. 

EXAMPLE  2.     If  10   liters  of  hydrogen  at  15°  C.  are 
burned,  what  volume  of  steam  at  300°  C.  is  formed  ? 
H2-}-0  =  H20. 

mam 

The  volume  of  the  steam  would  be  equal  to  that  of  the 
hydrogen,  if  the  temperatures  were  the  same,  making  the 
correction  for  the  change  of  temperature. 
10  (273  +  300)  =  5730  = 
273  +  15  288 

When  the  mass  of  a  solid  or  liquid  is  given  or  asked,  and 
the  volume  of  a  gas  is  asked  or  given,  the  equation  can 
only  be  solved  in  terms  of  the  mass  of  the  gas. 

EXAMPLE  3.  How  much  lead  sulphide  can  be  precipi- 
tated by  17  liters  of  hydrogen  sulphide? 


.^  HNO3  +  PbS. 
34  239 

17  liters  of  H2S  weigh  17*  34  grams. 

34  grams  of  H2S  precipitate  239  grams  of  lead  sulphide. 

239 
1  gram  of  H2S  precipitates  -  grams  of  PbS. 

17x  84  grams  of  H,S  precipitate 

17X84  x  —  =  181.3  grams  of  lead  sulphide. 


STOIC  HIOMETKY.  25 

53.  Gaseous  Diffusion.  Graham's  Law.  "The  veloc- 
ity of  the  diffusion  of  any  gas  is  inversely  proportional  to 
the  square  root  of  its  density."  This  law  applies  of  course 
to  volumes.  That  is,  when  two  gases  diffuse  through  the 
same  apparatus  for  equal  times  under  similar  conditions, 
the  volume  of  the  one  gas  diffused  multiplied  by  the 
square  root  of  its  density  is  equal  to  the  volume  of  the 
other  gas  diffused  multiplied  by  the  square  root  of  its 
density. 


EXAMPLE.  4  liters  of  hydrogen  diffuse  through  an 
apparatus  in  10  minutes,  and  1  liter  of  oxygen  in  an  equal 
time  under  similar  conditions  ;  find  the  density  of  oxygen. 


Z>=16. 


For  method  of  determining  the  empirical  formula  of  a 
substance  from  its  percentage  composition,  and  for  methods 
of  determining  the  relative  density  of  solids,  liquids,  and 
gases,  see  Appendix,  pp.  93  and  95. 


26  CHEMICAL   ARITHMETIC. 


EXAMPLES. 

MOLECULAR  WEIGHTS. 

1.  Find  the  molecular  weight  of  (a)  carbon  monoxide, 
CO;   (b)  magnesia,  MgO;    (c)  lime,  CaO;    (d)  alumina, 
A1203. 

2.  Find  the  percentage  of  oxygen  in  each  of  the  above- 
mentioned  bodies. 

3.  Find  the  molecular  weight  of  (a)  nitric  oxide,  NO; 
(£>)  sodium  hydrate,  NaHO ;  (c)  ferric  oxide,  Fe2O3. 

4.  Find   the    molecular   weight   of    (a)   zinc  sulphate, 
ZnSO4.7H2O;    (5)  copper  sulphate,  CuSO4.5H2O;    <V) 
sodium  sulphate,  Na2SO4.10H2O. 

THE  VOLUME  OF  A  MASS  or  GAS. 

1.  1000  cc.  of  gas  are  heated  from  0°  to  39°.     Find  the 
new  volume. 

2.  1000  cc.  of  gas   are   heated  from  39°  C.  to  52°  C. 
Find  the  new  volume. 

3.  The  pressure  on  10  cc.  of   gas  is  7  meters ;   if  the 
pressure  be  reduced  to  847  mm.,  what  is  the  new  volume? 

4.  1000  cc.  of  air  at  13°  C.  occupy  what  volume  at  65°  C.? 


EXAMPLES.  27 

5.  If  300  cc.  of  gas  are  measured  off  at  28°  C.,  what  will 
the  volume  become  at  — 14°  C.  ? 

6.  A  liter  of  gas  is  heated  from  14°  C.  to  42°  C.     Find 
the  new  volume. 

7.  A  liter  of  air  at  39°  C.  is  cooled  to  —  26°  C.    Find  the 
new  volume. 

8.  50  cc.    of  gas   at   10°  C.    occupy   what   volume   at 
24°  C.? 

9.  100  cc.  of  air  at  12°  C.  are  heated  until  they  occupy 
145  cc.     Find  the  new  temperature. 

10.  320  cc.  of  gas  are  measured  off  at  91°  C.  and  950  mm. ; 
what  is  the  normal  volume  ? 

11.  542  cc.  of  air  at  269°  C.  and  900  mm.  are  cooled  to 
51°  C.,  the  pressure  being  decreased  to  666  mm.    Find  the 
new  volume. 

12.  546  cc.  of  gas  at  17°  C.  and  760  mm.  are  cooled  to 
0°  C.,  the  pressure  being  decreased  to  600  mm.     Find  the 
new  volume. 

13.  A  quantity  of  oxygen,  which  measures  230  liters 
at  14°  C.  and  740  mm.,  will  measure  what  at  0°  C.  and 
760  mm.  ? 

14.  1234  cc.  of  normal  gas  are  cooled  to  —  52°  C.,  the 
pressure   being   decreased   to   617  mm.      Find   the   new 
volume. 


28  CHEMICAL  ARITHMETIC. 

MISCELLANEOUS  EXAMPLES. 

1.  50  grams  of  potassium  chlorate  are  heated;  what  mass 
of  oxygen  is  given  off  ? 

KC1O3  = 


2.    How  much  potassium  chlorate  is  required  to  make 
112  liters  of  oxygen  ? 

KC103  = 


3.    What  volume  of  oxygen  can  be  obtained  from  20 
grams  of  manganese  dioxide  by  heating  it  alone  ? 

2  =  Mn8O4-f  O2. 


4.   What  mass  of  oxygen  can  be  obtained  by  heating 
522  grams  of  manganese  dioxide? 

2  =  Mn3O4-f-O,. 


5.  What  mass  of  oxygen  (O2)  at  10°  C.  and  750  mm. 
would  fill  a  globe  of  15  liters  capacity  ? 

6.  On  heating  some  potassium  chlorate  298  grains  of 
potassium  chloride  were  left.     What  mass  of  chlorate  was 
heated,  and  what  mass  of  oxygen  was  formed  ? 

KC1O3  =  KC1  +  3O. 

7.  How  much  potassium  chlorate  is  required  to  make 
70  liters  of  oxygen  ? 

8.  60  grams  of  mercuric  oxide  are  heated  ;  what  volume 
of  oxygen  at  91°  C.  and  380  mm.  is  given  off? 


9.   270  grams  of  mercuric  oxide  are  heated  ;  what  volume 


EXAMPLES.  29 

of  oxygen  at  0°  C.  and  760  mm.  is  given  off?     And  what 
will  the  gas  measure  at  17°  C.  and  700  mm.  ? 

HgO  =  H 


10.  What  is  the  mass  of  13  liters  of  oxygen  (O2)  meas- 
ured at  12°  C.? 

11.  100  grams  of  steam  are  passed  over  red-hot  iron. 
What    volume    of   hydrogen    at    10°  C.  and   742   mm.    is 

formed? 

4  H2O  +  3  Fe  =  Fe3O4  +  4  H2. 

12.  How  many  liters  of  hydrogen  are  obtained  on  dis- 
solving 16  grams  of  magnesium  in  dilute  sulphuric  acid  ? 

Mg  +  H2S04  =  MgS04  +  H2. 

13.  What  volume   will  1000   cc.  of  hydrogen   at  0°  C. 
occupy  at  (a)  15°  C.,  (ft)  100°  C.,  (c~)  300°  C.? 

14.  What  volume  will  1000  cc.  of  hydrogen  at  100°  C. 
occupy,  at  (a)  0°  C.,  (ft)  -20°  C.,  (c)  -50°  C.? 

15.  4  liters  of  hydrogen  are  passed  over  heated  cupric 
oxide  ;  what  loss  of  mass  does  the  oxide  undergo  ? 


16.  Find  the  percentage  of  hydrogen  in  (#)  HC1,    (ft) 
H2S,  O)  NH4. 

17.  How  much  zinc  is  required  to  obtain  100  liters  oi 
hydrogen  at  91°  C.  and  800  mm.  ? 

Zn  +  H2SO4  =  ZnSO4  +  H2. 

18.  What  volume  of  oxygen  at  12°  C.  and  762  mm.  is 
evolved  on  decomposing  10  grams  of  water  by  electricity  ? 


30  CHEMICAL   ARITHMETIC. 

19.  500  cc.  of  hydrogen  at  39°  C.  are  exploded  with 
500  cc.  of  oxygen  under  a  pressure  of  332.5  mm.  ;  what 
volume  of  which  gas  is  left  ? 

=  H0. 


20.  2  grams  of  steam  are  passed  over  red-hot  iron  ;  what 
volume  of  hydrogen  at  10°  C.  and  770  mm.  is  formed? 

4  H2O  +  3  Fe  =  Fe3O4  +  4  H2. 

21.  100  grams  of  gypsum  are  heated;  what  volume  of 
steam  at  300°  C.  is  given  off? 

CaSO4,  2  H2O  =  CaSO4  +  2  H2O. 

22.  100   grams    of   each  variety    of   sodium    carbonate 
(a)  Na2CO3,  10H2O;     (5)  Na2CO3,  8H2O;     (c)   Na2CO3, 
5H2O;    (tT)  Na2CO3,  H2O,  contain   how  many  grams  of 
water  ? 

23.  What  do  100  liters  of  nitrogen  (N2)  weigh  ? 

24.  146  liters  of  nitrogen  at  17°  C.  and  974  mm.  are 
heated  to  51°  C.,  the  pressure  being  decreased  to  760  mm. 
Find  the  new  volume. 

25.  Find   the   percentage    composition    of    ammonium 
nitrate  NH4NO3. 

26.  How  many  pounds  of  nitric  acid  are  obtained  on 
distilling  400  pounds  of  sodium  nitrate  with   sulphuric 

acid? 

2  NaNO3  +  H2SO4  =  Na2SO4  +  2  HNO3. 

27.  How  much  copper  is  required  to  form  10  liters  of 
nitric  oxide  ? 

3  Cu  +  8  HN03  =  3  GuN  A  +  4  H2O  +  2  NO. 


EXAMPLES.  31 

28.  What  is  the  mass  of  270  cc.  of  nitrogen  measured 
over  water  at  8°  C.  and  768  mm.  ? 

29.  What  is  the  volume  of  20  grams  of  ammonia  (NH3) 
at  12°  C.  and  730  mm.  ? 

30.  25.82  liters  of  nitric  oxide  diffuse  through  a  certain 
apparatus  in  50  minutes  ;  what  volume  of  hydrogen  will 
diffuse  under  same  conditions? 

31.  The  pressure  on  134  cc.  of  air  is  increased   from 
480  mm.  to  1200  mm.     Find  the  new  volume. 

32.  What  volume  of  air  containing  21  per  cent  of  oxygen 
by  volume  is  required  to  burn  248  grams  of  phosphorus? 

P2  +  O5  =  P2O5. 

33.  How  much  phosphorus  is  required  to  remove  the 
oxygen  from  a  liter  of  air  ? 


34.  10  liters  of  carbon  monoxide  at  14°  C.  and  760  mm. 
are  required  ;  what  volume  of  normal  carbon  dioxide  must 
be  passed  over  red-hot  carbon,  and  what  mass  of  carbon  is 
absorbed  ? 


35.  A  substance  contains  carbon  20  per  cent,  oxygen 
26.6    per    cent,    and    sulphur   53.3    per    cent.      Find    its 
formula. 

36.  A  diamond  weighing  7  grams  is  burnt  in  oxygen; 
what  volume  of  carbon  dioxide  is  formed  ? 

2  =  C02. 


32  CHEMICAL   ARITHMETIC. 

37.  10  grams  of  turpentine  are  burnt;  what  volume  of 
carbon  dioxide  is  formed? 

C10H1(5  +  28  O  =  8  H2O  +  10  CO2. 

38.  How  much  marble  must  be  dissolved  in  acid  to  give 
20  liters  of  carbon  dioxide  at  18°  C.  and  740  mm.? 

CaCO3  +  2  HC1  =  CaCl2  +  H2O  +  C02. 

39.  How  much  hydrogen  potassium  carbonate  is  required 
to  give  a  liter  of  carbon  dioxide  ? 

HKCO3  +  HC1  =  KC1  +  H2O  +  CO2. 

40.  What  volume  is  occupied  by  177.5  grams  of  chlorine 


41.  How  much  salt  is  required  to   make  28  liters  of 
chlorine  ? 

2  NaCl  +  MnO2  +  2  H2SO4  =  2  H2O  +  Na2SO4  +  MuSO4  +  C12. 

42.  500  grams  of  potassium  chlorate  are  heated  ;  how 
much  potassium  chloride  is  left? 

43.  Find  the  percentage  composition  of  bleaching  pow- 
der, if  its  formula  be  CaO2Cl2. 

44.  From  2078  grams  of  sodium  chloride  what  volume 
of  chlorine  can  be  obtained  ? 

2  NaCl  +  MnO8  +  2  H2SO4  =  2  H2O  +  Na2SO4  +  MnS04  +  C12. 

45.  Find  the  percentage  composition  of  (a)  galena,  PbS; 
(6)  zinc  blende,  ZnS. 

46.  9.6  grams  of  sulphur  are  heated  in  hydrogen  ;  what 
volume  of  hydrogen  sulphide  is  formed  ? 

Jb  -f*  H2  =  H2o. 


EXAMPLES.  33 

47.  16  liters  of  hydrogen  diffuse  through  an  apparatus 
in  100  minutes;  how  much  sulphur  dioxide  (SO2)  will  dif- 
fuse under  the  same  conditions  ? 

48.  What  volume  of  sulphur  dioxide  at  20°  C.  and  740 
mm.  can   be   obtained  by  the  action  of  20  grams  of  sul- 
phuric acid  upon  copper  ? 

Cu  +  2  H2SO4  =  CuSO4  +  2  H2O  +  SO2. 

49.  100  grams  of  lead  form  146.45  grams  of  lead  sul- 
phate ;  find  the  molecular  weight  of  sulphuric  acid. 

Pb  +  H2SO4  =  H2  +  PbSO4. 

50.  The  skeleton  of  a  man  weighs  24  Ibs.,  and  contains 
58  per  cent  of  calcium  phosphate  (Ca32  PO4).      Find  the 
quantity  of  phosphorus  present. 

51.  How  much  phosphorus  can  be  obtained  from  169 
tons  of  bones  containing  53.7  per  cent  of  calcium  phos- 
phate (Ca3.2PO4)? 

52.  Find   the   percentage    composition    of  crystallized 
hydrogen  sodium  phosphate  (Na2HPO4. 121^0). 

53.  How  many  grams  do  10  liters  of  hydrogen  phos- 
phide weigh? 

54.  10  liters  of  hydrogen  diffuse  through  a  certain  appa- 
ratus in  a  certain  time  ;  what  volume  of  hydrogen  phosphide 
(PH3)  will  diffuse  under  similar  conditions? 

55.  In  a  case  of  poisoning,  11.73  grains  of  arsenic  trisul- 
phide  were  found ;  to  how  much  arsenic  trioxide  does  this 
correspond  ? 

As2S3  +  9  O  =  As  A  +  3  SO2. 


34  CHEMICAL   ARITHMETIC. 

56.  A  room  15  feet  long  and  10  feet  wide  and  high  is 
covered  with  a  paper  containing  .78  grains  of  Scheele's 
green  (CuHAsO3)  per  square  foot.     How  much  arsenic  is 
there  in  the  room? 

57.  To  make  a  kilogram  of   potassium    hydrate,  how 
much  («)  potassium  carbonate  and  (£)  calcium  hydrate  is 
required  ? 

K2C03  +  CaH202  ==  CaC03  +  2  KHO. 

58.  How  much  dry  sodium  carbonate  can  be  made  from 
500  kilograms  of  common  salt  ? 

2  NaCl  +  H2S04  =  2  HC1  +  Na2SO4. 

Na2S04  +  CaCO3  +  4  C  =  4  CO  +  CaS  +  NaXO3. 


PART    II. 

ELEMENTARY   QUALITATIVE   ANALYSIS. 
SECTION  I. 

PEELIMINARY   EXAMINATION   OF    SINGLE   SALTS   CONTAIN- 
ING  ONE   ACID   AND   ONE   BASE. 

TABLE  I.     Examination  for  Acid.* 

54.  If  the  substance  is  a  solid,  powder  a  portion  of  it, 
and  heat  in  a  glass  tube  sealed  at  one  end.  Notice  whether 
water  is  given  off.  If  so,  test  its  action  with  litmus  paper. 

Acid  reaction  indicates  sulphites,  chlorides,  etc. 

Alkaline  reaction  indicates  salts  of  ammonium  (NH4). 

If  a  sublimate  forms :  yellow  (or  in  red  globules)  indi- 
cates sulphur;  white>  ammonium  salts,  mercury,  arsenic, 
antimony. 

Metalic  mirror  indicates  arsenic. 

If  a  gas  is  evolved :  oxygen  indicates  chlorates,  nitrates, 
peroxides;  carbon  monoxide  indicates  oxalates ;  nitrogen 
tetroxide  indicates  nitrates  ;  ammonia  indicates  ammonium 
salts  ;  carbon  dioxide  indicates  carbonates. 

If  the  substance  alters  in  color :  black  indicates  organic 
matter ;  yellow  while  hot  indicates  zinc  oxide. 

*  Mercury,  sulphur,  ammonia,  though  not  acids,  arc  included. 


36  ELEMENTARY    QUALITATIVE   ANALYSIS. 

Take  another  portion  of  substance  under  analysis,  and 
add  HC1.  Notice  whether  a  gas  is  evolved  with  efferves- 
cence. 

If  it  smell. like  burning  sulphur,  it  indicates  sulphites, 
or  hyposulphites. 

If  it  has  the  odor  of  rotten  eggs,  sulphides. 

If  it  has  the  odor  of  bitter  almonds,  cyanides. 

If  it  has  the  odor  of  chlorine  on  heating,  peroxides, 
chromates,  or  hypochlorites. 

If  it  renders  lime-water  turbid,  carbonates. 

Take  another  portion  of  substance,  and  try  if  it  is 
soluble  in  water ;  if  so,  add  BaCl2  solution  to  a  portion  of 
the  solution  and  notice  whether  a  precipitate  form. 

A  white  precipitate  insoluble  in  HC1  indicates  sulphates. 

White  and  soluble  in  HC1  indicates  phosphates,  silicates, 
oxalates,  borates,  and  fluorides,  also  carbonates  and  sulphites. 

Yellow  indicates  chromates. 

If  BaCl2  gives  no  precipitate,  add  AgNO3  to  another 
portion  of  the  solution  and  notice  if  a  precipitate  form. 
White  precipitate  indicates  chlorides,  also  cyanides. 
Yellowish-white  indicates  bromides  and  iodides. 
Black  indicates  sulphides. 

In  case  neither  water  nor  HC1  has  dissolved  the  sub- 
stance, try  HNO3. 

If  this  does  not  dissolve  it,  try  aqua  regia ;  and  if  that 
fails,  try  method  described  in  Table  II. 

If  the  substance  is  dissolved  in  HNO3  or  aqua  regia,  it 
must  be  evaporated  to  dry  ness  with  HC1  before  proceed- 
ing to  the  examination  for  base. 


EXAMINATION   FOR   BASE.  37 


SECTION   II. 
Examination  for  Base. 

55.  Having  obtained  a  solution,  add  HC1.  If  it  pro- 
duces a  precipitate,  it  indicates  silver,  had,  or  mercurous 
salts.  Add  HC1-|-H2S.  If  it  produces  a  precipitate, 

Black  indicates  mercuric  salts,  lead,  bismuth,  or  copper  ; 

Yellow  indicates  arsenic,  stannic  salts,  or  cadmium; 

Orange  indicates  antimony  ; 

Brown  indicates  stannous  salts. 

If  (NH4)HO  +  (NH4)C1+(NH4)2S  produce  a  precipi- 
tate, it  indicates 

Black,  iron,  nickel,  cobalt; 
White,  zinc  or  aluminum  ; 
Flesh-colored,  manganese; 
Green,  chromium. 

If  (NH4)HO  +  (NH4)C1+(NH4)2CO3  produce  a  precipi- 
tate, it  indicates 
\     Barium  (tinges  flame  green), 

Strontium  (tinges  flame  crimson), 
Calcium  (tinges  flame  dull  red). 

If  the  solution  is  not  precipitated  by  any  of  the  above 
reagents,  it  indicates  magnesium,  potassium,  sodium,  am- 
monium, of  which  the  following  are  the  individual  tests:  — 

Magnesium  is  precipitated  by  Na2HPO4  h  (NH4)HO, 
white. 

Potassium  is  precipitated  (except  in  very^jlilute  solu- 

'' 


I17SESIT7 


38  ELEMENTARY   QUALITATIVE   ANALYSIS. 

tions)  by  PtCl4,  precipitate  insoluble  in  alcohol ;  also 
tinges  the  flame  violet. 

Sodium  is  precipitated  by  H2SiF6 ;  also  tinges  flame  in- 
tense yellow,  not  visible  through  blue  glass. 

Ammonium  salts  heated  with  NaHO  give  smell  of  NHa. 


TABLE  II. 

Examination  of  Insoluble  Substances.  The  follow- 
ing substances  are,  under  certain  circumstances,  insoluble 
in  acids,  and  must  be  examined  specially:  — 

Silica,  /Silicates. 

Alumina,  Aluminates. 

Oxides  of  Antimony,  Chromium,  and  Tin. 

Chrome  Iron  Ore. 

Sulphates  of  Barium,  Strontium,  and  Lead. 

Certain  Fluorides  (e.g.  of  Calcium). 

Certain  Sulphides  (e.g.  of  Lead). 

Chloride,  Bromide,  and  Iodide  of  Silver. 

Carbon. 

Sulphur. 

Heat  the  substance  in  a  dry  tube  as  before,  and  notice 
if  it  fuses  and  volatilizes  completely.  If  it  smells  of  SO2, 
it  indicates  sulphur. 

If  it  fuses,  but  does  not  volatilize,  indicates  chloride, 
bromide,  or  iodide  of  silver  (also  will  yield  metallic  silver 
on  fusing  on  charcoal  with  Na2CO3). 

If  it  is  infusible,  but  disappears  on  heating,  carbon  (de- 
flagrates when  heated  with  KN03). 

If  it  is  infusible,  but  darkened  in  color  while  hot,  regain- 
ing its  color  on  cooling,  tin  dioxide  and  antimony  pentoxide 


' 


EXAMINATION    OF   INSOLUBLE   SUBSTANCES.  39 

(confirmed  by  blow-pipe  test  —  tin  bead  malleable ;  anti- 
mony bead,  brittle). 

Notice  whether.  It  yields  a  green  bead  with  borax  or 
microcosmic  salt ;  it  indicates  chromium  oxide,  or  chrome 
iron  ore.  It  swims  undissolved  in  a  bead  of  microcosmic 
salt,  silica  and  silicates  (fuse  with  four  times  its  weight  of 
a  mixture  of  K2CO3  and  Na2CO3.  Allow  to  cool,  dissolve 
in  water,  add  HC1,  and  evaporate  to  dryness.  Silica  will 
separate  out  as  a  gelatinous  mass). 

It  yields  a  colorless  bead,'  with  microcosmic  salt,  alu- 
mina. (Heated  on  charcoal,  and  moistened  with  CO(NO3)2 
and  reheated,  it  yields  a  blue,  infusible  mass.) 

It  is  white  and  infusible,  but  quite  unaltered  by  heat- 
ing. 

Lead  sulphate  yields,  when  heated  with  Na2CO3  in  blow- 
pipe reducing  flame,  malleable  metallic  bead. 

Barium  sulphate  fused  with  Na2CO3  yields  BaCO3.  Boil 
the  fused  mass  with  water,  filter  and  wash;  the  residue 
dissolved  in  HC1  yields  BaCl2  (flame  color  green),  precipi- 
tated by  SrSO4  solution. 

Strontium  sulphate  fused  with  Na2CO3  yields  SrCO3.  Boil 
the  fused  mass  with  water,  filter  and  wash;  the  residue 
dissolved  in  HC1  yields  SrCl2  (flame  color  crimson),  preci- 
pitated by  CaSO4  solution. 

Calcium  fluoride  heated  with  H2SO4  yields  HF,  which 
etches  glass. 

It  is  black  and  infusible,  and  yields  a  malleable  metallic 
bead  when  fused  with  Na2CO3  in  the  blow-pipe  flame. 

Lead  sulphide  (bead  leaves  mark  on  paper),  and  when 
dissolved  in  HNO3  gives  a  white  precipitate  on  addition 
of  H2SO4. 

The  action  of  strong  H2SO4  often  affords  a  valuable  indi- 


40  ELEMENTARY   QUALITATIVE   ANALYSIS. 

cation  of  the  nature  of  a  salt,  whether  soluble  or  in- 
soluble. ' 

Thus  evolution  of 

Sulphur  dioxide  indicates  sulphites  or  hyposulphites. 

Sulphuretted  hydrogen  indicates  sulphides. 

Hydrocyanic  acid  indicates  cyanides. 

Oxygen  indicates  peroxides,  chromates,  permanganates. 

Carbon  dioxide  indicates  carbonates. 

Carbon  monoxide  indicates  oxalates,  formates,  ferro* 
cyanides. 

Chlorine  indicates  hypochlorites. 

Hydrochloric  acid  indicates  chlorides. 

Hydrofluoric  acid  indicates  fluorides. 

Nitric  acid  indicates  nitrates. 

Acetic  acid  indicates  acetates. 

Chlorine  tetroxide  indicates  chlorates. 


REACTIONS  OF  THE  COMMONLY  OCCURRING  METALS  WITH 
THE  METHODS  OF  SEPARATION. 

GROUPING  OF  THE  METALS. 

56.  The  metals  are  divided  into  five  groups,  according 
to  their  behavior  with  certain  substances  termed  group 
reagents. 

Group  I.     (Silver  Group.) 

Group  reagent,  HC1.  Metals  whose  chlorides  are  insolu- 
ble in  water.  They  are  precipitated  from  the  solutions  of 
their  salts  by  the  first  group  reagent,  hydrochloric  acid. 

Silver,  mercury  (rnercurous  salts),  lead. 


GROUPING  OF   THE  METALS.  41 

Group  II.     (Copper  Group.) 

Group  reagent  H2S  in  presence  of  HC1.  Metals  which 
in  acid  solutions  form  insoluble  sulphides,  are  precipitated 
from  their  acidulated  solutions  by  the  second  group  rea- 
gent H2S  (hydrosulphuric  acid). 

Arsenic,  antimony,  tin,  lead,  bismuth,  copper,  cadmium, 
mercury  (mercuric  salts). 

The  three  metals,  arsenic,  antimony,  and  tin,  form  a 
sub-group,  as  their  sulphides  are  soluble  in  (NH4)2S2,  whilst 
the  sulphides  of  the  remaining  metals  are  insoluble  in  that 
reagent. 

Group  III.     (Iron  Group.) 

Group  reagent  (NH4)2S  in  presence  of  (NII4)C1  and 
(NH4)HO. 

Metals  whose  sulphides  and  hydroxides  are  insoluble  in 
water,  but  decomposed  by  dilute  acids,  are  precipitated 
from  neutral  solutions  by  the  third  group  reagent,  ammo- 
nium sulphide.  Aluminium,  and  chromium  are  precipi- 
tated as  hydrates ;  the  others  as  sulphides.  Iron,  nickel, 
cobalt,  zinc,  manganese,  as  sulphides. 

Group  IV.     (Barium  Group.) 

Group  reagent  (NH4)2CO3  in  presence  of  (NH4)HO  and 
(NH4)C1. 

Metals  whose  carbonates  are  insoluble  in  water,  and  are 
precipitated  from  their  solutions  by  the  fourth  group  rea- 
gent, ammonium  carbonate ;  barium,  strontium,  calcium. 

Group  V.     (Potassium  Group.) 

Metals  not  precipitated  by  any  of  the  above  group  rea- 
gents, as  their  chlorides,  sulphides  and  carbonates  are 


42  ELEMENTARY  QUALITATIVE  ANALYSIS. 

soluble  in  water.  They  are,  therefore,  distinguished  by 
individual  tests :  magnesium,  potassium,  sodium,  ammo- 
nium. 

57.  Each  group  reagent  will  precipitate  the  metals  of 
preceding  groups.  The  metals  distinguished  by  being 
insoluble  as  chlorides  (Group  I.)  are  also  insoluble  as 
sulphides  (with  Groups  II.  and  III.)  and  as  carbonates 
(with  Group  IV.).  The  second  group  sulphides  are  pre- 
cipitated both  from  acid  and  from  neutral  solutions,  though 
the  third  group  sulphides  are  precipitated  from  neutral, 
but  not  from  acid  solutions,  and  second  arid  third  group 
metals  form  insoluble  carbonates,  as  well  as  those  of 
Group  IV. 

In  the  work  of  analysis,  the  first  group  metals  may  be 
worked  with  the  second,  but  thereafter  the  metals  found 
in  each  group  must  be  completely  removed  before  testing 
for  the  next  group. 

After  filtering  out  a  group  precipitate,  the  reagent 
which  produced  it  should  be  again  carefully  applied,  with 
the  proper  conditions,  to  the  filtrate  before  testing  it  for 
the  next  group. 

The  student  should  at  first  have  several  metallic  salts 
given  to  him,  and  be  asked  merely  to  determine  to  which 
of  the  above  groups  each  salt  belongs. 

He  ought  next  to  make  himself  familiar  with  the  indi- 
vidual tests  for  each  metal  which  follows,  and  then  proceed 
to  the  separations  of  the  different  metals.  It  will  also  be 
well  for  him  to  attempt  to  frame  a  table  of  separations  for 
each  group  before  consulting  those  given  in  the  book. 


REACTIONS  OF  THE  METALS  OF  THE  SILVER  GROUP.     43 

r 

58.    Reactions  of  the  Metals  of  the  Silver  Group 

(Group  I.). 

Silver,  Ag',  108.     Solution  for  Reactions,  AgNO3. 

1.  HC1  produces  a  white,  curdy  .precipitate  of  AgCl, 
insoluble  in  hot  water,  soluble  in  Rff4fi<J  and  in  KCN ; 
reprecipitated  by  HNO3;  darkens  on  exposure  to  light. 

2.  H2S    or    (NH4)2S    produces   a    black   precipitate    of 
Ag2S,  .Soluble  in  boiling  HNO3,  with  separation  of  sulphur. 

3.  NailO  PSPdu/cg8  a  light-brown  precipitate  of  Ag2O, 
soluble  iii.  (NH45FOJ. 

4.  (KffJjHO   produces    (from  neutral  solutions   only) 
brown  Ag2O,  soluble  in  excess  of  reagent. 

5.  K2CrO4  produces  a  dark-red  precipitate  of  Ag2CrO4, 
soluble  in  hot  HNO3  or  in  (NH4)HO;  this  solution  de- 
posits   on    cooling    an   acid   chromate    in    needle-shaped 
crystals. 

6.  KI  precipitates  yellowish  Agl. 

7.  Cu  and  some  other  metals  precipitate  metallic  Ag. 

8.  Na2HPO4  precipitates  yellow  Ag3PO4. 

9.  Heated  on  charcoal  with  Na2CO3,  in  the  reducing 
flame  of  the  blow-pipe,  yields  bright,  malleable  metallic 
beads,  soluble  in  HNO3. 

Characteristic  reaction,  1. 

Mercury,  Hg",  200  (Mercurous  Salts).     Solution  for 
Reactions,  Hg2N2O6. 

1.  HC1  produces  white  precipitate  of  Hg2Cl2  (calomel), 
insoluble  in  cold  HNO3;  blackened  by   (NH4)HO,  from 
formation  of  Hg2Cl(NH2). 

2.  H2S    or    (NH4)2S    produces   a    black   precipitate   of 
Hg2S,  not  dissolved  by  boiling  HNO3. 


44  ELEMENTARY  QUALITATIVE  ANALYSIS. 

3.  NaHO  precipitates  black  Hg2O,  insoluble  in  excess 
of  NaHO  or   (NH4)HO ;    decomposes   readily  into   HgO 
and  Hg. 

4.  SnCl2  precipitates  gray  Hg.     If  the  fluid  be  poured 
off  and  the  residue  boiled  with  HC1,  distinct  globules  are 

.  obtained. 

5.  KI  precipitates  dark-green  IIg2L. 

6.  A  drop  of  a  metal  or  only  slightly  acid  solution  of 
a  mercurous    salt   placed    on    a   bright    copper   coin  will 
deposit   mercury,   and  the-  stain  will   become   bright  by 
rubbing. 

7.  Heated  in    small    tube   with    NaHCO3,   yields   gray 
deposit  of  Hg.    Hg  is  volatile,  and  condenses  on  the  cooler 
parts  of  the  tube';'  soluble  in  HNO3. 

Characteristic  reactions,  1,  7. 


Lead,  Pb",  207.     Solution  for  Reaction,  PbN2O6. 

1.  HC1  precipitates  (incompletely)  white  PbCl2,  solu- 
ble in  boiling  H2O,  or  in  large  quantity  of  cold  H2O ;  con- 
verted into  a  basic  salt  on  adding    (NEk)HO,   without 
change  of  appearance.     If  PbCl2  be  dissolved  in  boiling 
H2O,  it  will  crystallize  from  this  solution  on  cooling. 

2.  H2SO4   precipitates  heavy  white   PbSO4,  soluble  in 
NaHO    or    ammonium    tartarate.       This    precipitate    in 
dilute  solutions  only  appears  on  standing.     If  there  is  no 
immediate     precipitation,     concentrate    the    solution    by 
evaporation.     PbSO4  is   soluble  in  boiling  HC1,  and  the 
solution,  -on    cooling,    deposits    iieedle-shaped   crystals  of 
PbCl2. 

3.  H2S  or'(NH4;)2S  precipitates  black  PbS,  soluble  in 
hot  HNQ,. 


. 


REACTIONS  OF  THE  METALS  OF  THE  SILVER  GROUP.      45 

4.  K2CrO4  precipitates   bright-yellow    (chrome-yellow) 
PbCrO4;  soluble  in  NaHO ;  soluble  with  difficulty  in  HNO3. 

5.  KI  precipitates  bright-yellow  PbI2,  soluble  in  boil- 
ing H2O ;    the    solution  on  cooling   deposits  the   salt  in 
brilliant,  golden  hexagonal  crystals,  f 

6.  Zn   precipitates   metallic    Pb   in    crystalline    form. 
Known  as  the  "  lead  tree." 

7.  Heated  on  charcoal  with  NaHCO3,  yields  malleable 
beads  of  Pb,  soluble  in  HNO3,  and  at  the  same  time  a 
yellow  incrustation  of  PbO  on  the  charcoal. 

Characteristic  reactions,  2,  4,  5. 

59.    Table  II.  for  the  Separation  of  Silver  (Group  I.). 
Silver,  Mercury,  and  Lead. 

(#)  Add  HC1  until  no  further  precipitation  takes  place. 
Filter  from  the  precipitated  chlorides. 


Precipitate  contains 
AgCl,  Hg2Cl2,  PbCl2. 


Filtrate  contains 
Group  II.,  III.,  IV.,  and  V. 


(6)  Wash  precipitate  twice  with  cold  H2O,  and  add 
washings  to  nitrate;  then  twice  with  hot  H2O.  Test  part 
of  this  for  Pb*with  dilute  H2SO4.  White  precipitate  indi- 
cates lead.  Boil  remainder  down  to  obtain  crystals  of 
PbCl2.  If  Pb  is  found,  the  precipitate  is  washed  repeat- 
edly with  hot  H2O,  till  free  from  it.  Residue  indicates 
AgCl,  Hg2Cl2,  insoluble  in  hot  H2O. 

0)  To  residue  add  warm  (NIl^HG;  filter. 


Residue  is  Hg. 

.If  residue  is  black,  it  indi- 
cates mercury.  Dissolve  in 
HCl-fHNO3,  and  test  with 
SnCl2. 


Filtrate  Ag. 

Add  HNO3;  white  precipitate 
indicates  the  presence  of  silver. 


ELEMENTARY   QUALITATIVE   ANALYSIS. 
Reactions  of  the  Metals  of  the  Copper  Group. 

6O.  Metals  whose  sulphides  are  insoluble  in  HC1  and 
are  precipitated  in  presence  of  that  acid  by  the  group 
reagent  H2S. 

MERCURY,   LEAD,    BISMUTH,    COPPER,    CADMIUM,    ARSENIC, 
ANTIMONY,    AND   TIN. 

Sub-Group  A.  Sulphides  of  the  above  metals  insoluble 
in  (NH4)2S2?  viz. :  Mercury,  Lead,  Copper,  Bismuth,  and 
Cadmium. 

Mercury,  Hg"   (Mercuric  Salts).     Solution  for  Reactions, 

HgCl2. 

1.  H2S  produces,  when  added  by  degrees,  first  a  white 
precipitate,  which  changes  to  ofange,  then  to  brownish-red, 
and  finally  to  a  black  precipitate  of  HgS.    These  successive 
changes  of  color  on  the  addition  of  H2S  are  exceedingly 
characteristic.     This  precipitate  is  insoluble  in  HC1*  and 
in  HNO3,  even  on  boiling ;  it  i^sqlubl^  in  KHS  and  in 
aqua  regia. 

2.  KHO  produces  a  yellow  precipitate  of  HgO  insoluble 
m  excess,  except  when  added  to  very  acid  solutions. 

3.  (NH4)HO  produces  a  white  precipitate  of  HgCl(NH2) 
(white  precipitate). 

4.  SnCl2  when   added  in  small  quantities  precipitates 
Avhite  Hg2Cl2;   when  added  in  excess,  gray  metallic  Hg, 
which  may  be  united  into  a  globule  by  boiling  with  HC1. 

5.  KI  precipitates  bright-red  HgI2,  soluble  in  excess  of 
either  KI  or  HgCl2. 

6.  K2CrO4  precipitates  an  orange  basic  chromate  easily 
soluble  in  HNO,. 


REACTIONS  OF  THE  METALS  OF  THE  COPPER  GROUP.     47 

7.    Reactions  6  and  7  for  mercurous  salts  are  also  pro- 
duced with  mercuric. 

Characteristic  reactions,  1,  4. 

Lead,  Pb",  207.     Solution  for  Reactions,  PbN2O6. 

1.  H2S    precipitates   black   PbS,   even   in  solutions  of 
PbCl2,  so  that  lead  belongs  to  both  the   silver   and   the 
copper  groups. 

2.  Reactions  2,  3,  4,  5,  for  lead  in  Group  I.  are  also 
applicable  in  this  group. 

Bismuth,  Bi'",  210.     Solution  for  Reactions,  BiCl3. 

1.  H2S  precipitates  black  Bi2S3,  insoluble  in  KHS  and 
KHO,  but  soluble  in  HNO3. 

2.  KHO   or    (NH4)HO   precipitates    white    BiO.OH, 
which  on  boiling  becomes  yellow  (Bi2O3) ;  precipitate  is 
insoluble  in  excess  of  either  reag^ai. 

3.  H2O,  when  added  in  considerable  quantity  to  normal 
salts  of  bismutb,  precipitates  white  basic  salt  of  bismuth, 
BiOCl,  insoluble  in  tartaric  acid.     Solutions  of   bismuth 
salts  containing  much  free  acid  do  not  give  this  reaction 
with  H2O  until  the  excess  of  acid  has  been  expelled  by 
evaporation. 

4.  Zn  or  Fe  precipitates  spongy  Bi. 

5.  K2CrO4  precipitates  yellow  Bi2.3(CrO4),  soluble  in 
HNO3,  and  insoluble  in  NaHO. 

6.  Heated  on  charcoal  with  NaHCO3,  in  reducing  flame 
of  blow-pipe,    compounds    of    Bi   yield    brittle    metallic 
globules;  also  a  yellow  incrustation  of  Bi2O3  on  the  char-  - 
coal.     Bi  is  soluble  in  HNO3  or  aqua  regia. 

Characteristic  reactions,  3,  6. 


48  ELEMENT  AJRY   QUALITATIVE   ANALYSIS. 

Copper,  Cu",  63.5.     Solution  for  Reactions,  CuSO4. 

1.  H2S  precipitates  black,  CuS,  soluble  in  HNO3;  insolu- 
ble in  KHS,  and  only  slightly  soluble  in  (NH4)2S2.     CuS 
is  also  soluble  in  KCN,  but  insoluble  in  hot  dilute  H2SO4. 

2.  KHO  precipitates  a  pale-blue  Cu(HO)2,  insoluble  in 
excess.     If  KHO  be  added  in   excess   and   the   mixture 
boiled,  the  precipitate  becomes  black. 

3.  (NH4)HO  precipitates,  when  added  in  small  quanti-' 
ties,  greenish-blue  basic  salt,  soluble  in  excess  of  (NH4)HO, 
forming  a  dark-blue  solution   which   consists   of   double 
basic  salt  of  copper  and  ammonium. 

4.  K4Fe(CN)6  precipitates  brown  Cu2Fe(CN)6,  insoluble 
in  dilute  acids,  but  decomposed  by  KHO. 

5.  Fe  precipitates  Cu  in  the  metallic  state,  especially  in 
the  presence  of  a  little  free  acid.     > 

6.  Zn  also  precipitates  copper  solutions. 

7.  K2CrO4  precipitates  a  brownish-red  basic  chromate, 
soluble  in  HNO3  and  in  (NH4)H0. 

8.  Compounds  of  Cu,  when  heated  in  Bunsen  flame,  im- 
part a  green  color,  especially  after  addition  of  AgCl. 

9.  Heated  on  charcoal  with  NaHCO3  in  reducing  flame, 
yields  brittle  metallic  globules  of  bright-red  color,  soluble 
in  HNO3  or  concentrated  H2SO4. 

Characteristic  reactions,  3,  5,  6. 

Cadmium,  Cd",  112.     Solution  for  Reactions,  CdN2O6. 

1.  H2S  precipitates  yellow  CdS,  soluble  in  HNO3,  in- 
soluble in  KHS,  KCN,  and  (NH4)2S.     CdS  is  dissolved  by 
hot  dilute  H2SO4. 

2.  KHO  precipitates  Cd(HO)2,  insoluble  in  excess  of 
reagent. 


REACTIONS  OF  THE  METALS  OF  THE  COPPER  GROUP.     49 

3.  (NH4)HO  precipitates  Cd(HO)2;  soluble  in  excess 
of  reagent. 

4.  Zii  precipitates  Cd  in  brilliant  scales. 

5.  Heated  on  charcoal  with  NaHCO3  in  the  reducing 
flame,  yields  a  brown  incrustation  of  CdO.     Cd  dissolves 
readily  in  HNO3. 

Characteristic  reactions,  1,  5. 


50 


ELEMENTARY  QUALITATIVE   ANALYSIS. 


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REACTIONS  OF  THE  METALS  OF  THE  C 


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52  ELEMENTARY  QUALITATIVE  ANALYSIS. 

62.    Sub-Group  B.     Sulphides  soluble  in  (NH4)2S2,  viz.: 
those  of  Tin,  Antimony,  and  Arsenic. 

Tin,  Sn",  118  {Stannous  Salts).     Solution  for  Reactions, 

SnCl2. 

1.  H2S  precipitates  dark-brown  SnS,  soluble  in  yellow 
(NH4)2S2    (nearly   insoluble    in    colorless    (NH4)2S) ;    re- 
precipitated  as  yellow  SnS2  by  HC1. 

2.  KHO   precipitates   Sn(HO)2,    soluble   in    excess   of 
reagent. 

3.  (NH4)HO  precipitates  Sn(HO)2,  not  soluble  in  excess. 

4.  HgCl2  precipitates  at  first  white  Hg2Cl2;  on  boiling 
with  excess  of  reagent,  gray  Hg. 

5.  AuCL.,  on  addition  of  a  little  HNO8,  precipitates  pur- 
ple (royal  purple  of  Cassius). 

6.  Zn  precipitates  metallic  tin. 

7.  Mixed  with  NaHCO3-f  KCN,  and  heated  on  charcoal 
in  reducing  flame,  yields  small  globules  of  Sn  and  a  white 
incrustation  of  SnO2. 

Stannic  Salts,  Sniv.     Solution  for  Reactions,  SnCl4. 

Stannous  salts  are  converted  into  stannic  by  oxidizing 
agents,  HNO3,  Cl,  Fe2Cl6,  etc. 

1.  H2S  precipitates  yellow  SnS2,  soluble  in  (NH4)2S,  in 
KHO,  and  in  boiling  concentrated  HC1.     It  is  with  diffi- 
culty soluble  in  (NH4)HO,  and  insoluble  in  (NH4)2CO3. 

2.  KHO  or  (NH4)HO  precipitates  white  SNO(HO)2, 
soluble  in  an  excess  of  precipitant. 

3.  Zn  produces  same  reaction  as  with  stannous  salts. 

4.  The  blow-pipe  reaction  for  stannic  is  the  same  as  for 
stannous. 

Characteristic  reaction  :  stannous,  1,  4 ;  stannic,  1. 


REACTIONS  OF  THE  METALS  OF  THE  COPPER  GROUP.   53 

63.  Antimony,  Sb'^,'  122.3.    Solution  for  Reactions,  SbCl3. 

1.  H2S  precipitates  orange   Sb2S3,    soluble  in  (NH4)2S, 
NaHO,    and    in    hot    concentrated    HC1 ;     insoluble    in 
(NH4)2C03. 

2.  KHO  precipitates  Sb2O3,  soluble  in  excess  of  reagent. 

3.  (NH4)HO  precipitates  Sb2O3,  insoluble  in  excess. 

4.  H2O  produces  in  solutions  of  SbCl3  a  white  precipi- 
tate of  SbOCl,  soluble  in  tartaric  acid. 

5.  Zn  in  presence  of  HC1  and  platinum  precipitates  Sb 
as  a  black  powder,  which  adheres  to  the  platinum.     The 
black  stain  on  the  platinum  is  not  removed  by  HC1,  but 
is  immediately  dissolved  by  warm  HNO3. 

6.  (Marsh's  Test.)     If  a  solution  of  Sb  be  placed  in  a 
flask  in  which  hydrogen  is  being  generated,  SbH3  is  given 
off  as  a  gas,  which  is  decomposed  by  heat,  Sb  being  de- 
posited.    This  is  best  done  by  holding  in  the  SbH3  flame 
a  piece  of  cold  porcelain  ;  a  dull-black  stain  of  metallic  Sb 
will  be  deposited  on  it.     Add  to  the  stain  on  the  porcelain 
a  drop  of  NaCIO ;  the  stain  will  remain  undissolved. 

7.  Heated  with  NaHCO3  on  charcoal  in  the  reducing 
blow-pipe  flame,  yields  brittle  globules  of  the  metal  and  a 
white  incrustation  of  Sb2O3  on  the  charcoal. 

Characteristic  reactions,  5,  6. 

64.  Arsenic,  As'"  (Arsenious  Compounds).     Solution  for 

Reactions,  Na3AsO3. 

1.  H2S   (in  acid   solutions)   precipitates  yellow  As2S3, 
soluble  in  alkaline  sulphides  in  KHO,  in  HNO3,  and  in 
(NH4)2CO3,  but  nearly  insoluble  in  boiling  concentrated 
HC1. 

2.  (NH4)HO  and  NaHO  produce  no  precipitates. 


54  ELEMENTARY  QUALITATIVE  ANALYSIS. 

3.  AgNO3  produces  in  neutral  solutions  a  pale-yellow 
precipitate  of  Ag3AsO3. 

4.  CuSO4  added    to    a    neutral    solution    precipitates 
CuHAsOs  (Scheel's  green),  soluble  in  (NH4)HO. 

5.  Cu  added   to  an  HC1  solution  of  arsenic  becomes 
coated  with  a  gray  film  of  metallic  As.     Reinscfts  test. 

6.  Proceed  exactly  as  in  Marsh's  test  for  Sb,  substi- 
tuting a  solution  of  As  for  one  of  Sb,  and  observe  the 
bluish  flame  with  which  the  mixture  of  H  and  AsH3  burns, 
and  also  the  production  of  white  fumes,  As2O3.     Obtain 
stains  on  cold  porcelain  as  in  the  case  of  Sb.     Observe  the 
distinction  in  color  of  the  stains.     Dark-brown  or  almost 
black,  in  the  case  of  Sb ;  pale-brown  and  lustrous  in  the 
case  of  As.     Add  to  one  of  the  stains  on  porcelain  a  drop 
of  NaCIO  ;  it  will  be  rapidly  dissolved. 

7.  Arsenious    compounds   are   converted   into   arsenic 
compounds  by  oxidizing  agents. 

Arsenic  Compounds,  Asv.    Solution  for  Reactions,  Na3AsO4. 

1.  H2S  in  warm  solutions  free  from  HNO3  precipitates 
As2S3  and  S. 

2.  See  2,  under  Arsenious  Compounds. 

3.  AgNO3  .precipitates  from  neutral  solutions  reddish- 
brown  Ag3AsO4. 

4.  MgSO4  in  presence  of  (NH4)HO  and  (NH4)C1  pre- 
cipitates white  crystalline  MgNH4AsO4. 

5.  Heated  on  charcoal  with  Na2CO3,  all  compounds  of 
arsenic  are  reduced  to  As,  which  volatilizes  with  character- 
istic garlic  odor. 

6.  See  6,  under  Arsenious  Compounds,  Marsh's  test. 
Characteristic  tests :  arsenious,  4,  5,  6 ;  arsenic,  5,  6. 


REACTIONS  OF  THE  METALS  OF  THE  IRON  GROUP.      55 

TABLE   IV. 

Group  II. 

65.    Separation  of  Arsenic,  Antimony,  and  Tin  (Sub-Group  B). 

Solution  in  (NH4)2S2  contains  sulphides  of  As,  Sb,  and  Sn.  Add  HC1 
until  acid :  the  metals  are  re-precipitated  as  sulphides.  Filter ;  wash  pre- 
cipitate with  hot  water  till  free  from  HC1 ;  digest  precipitate  with  one  or 
two  pieces  of  solid  (NHJ2CO8  and  H20.  Filter. 

RESIDUE.  FILTRATE. 

SnS2,  Sb2S3.  As. 

Wash    and    dissolve    in    strong  Add  HC1  until  acid;  wash  pre- 

boiling  HC1,  dilute  with  water,  and  cipitated  sulphide,  and  dissolve  in 

add  to  the  solution  a  piece  of  zinc  HC1  and  a  little  KC103,  boil  down 

and  platinum  :    Sb  forms   a  black  to  a  small  bulk,  and  apply  Marsh's 

stain  on  the  platinum.     Dissolve  in  test.       Metallic     mirror,     yielding 

HNO3,  boil  down  to  a  small  bulk,  octahedral     crystals     on     heating, 

and  pass  H2S  through  the  solution :  indicates     Arsenic.       Dissolve     in 

orange    precipitate   indicates   pres-  H2O,  and  confirm  by  adding  AgNO3 

ence  of  Antimony.   The  Sn  deposits  and  dilute  (NH4)HO,  to  obtain  yel- 

on  the  zinc ;  dissolve  in  HC1,  boil  low  precipitate  of  Ag3As03. 
down,  and  test  with  HgCl2.     White 
precipitate  indicates  Tin. 

66.    Keactions  of  the  Metals  of  the  Iron  Group 

(Group  III.). 

Metals  whose  sulphides  and  hydrated  oxides  are  insolu- 
ble in  water,  and  are  precipitated  on  addition  of  the  group 
reagent  (NH4)2S  in  presence  of  (NH4)HO  and  (NH4)C1. 

IRON,   NICKEL,   COBALT,   ZINC,   ALUMINUM,   MANGANESE, 
AND   CHROMIUM. 

Iron,  Fe",  56  (Ferrous  Salts').  Solution  for  Reactions,  FeCl2. 

1.  (NH4)2S  precipitates  black  FeS,  soluble  in  HC1,  in- 
soluble in  alkalies. 

me 


56  ELEMENTARY  QUALITATIVE   ANALYSIS. 

2.  (NH4)HO   or   KHO   precipitates    white    Fe(HO)2, 
which  rapidly  acquires  a  dirty-green  color,  and  ultimately 
a  reddish-brown  color,  owing  to  absorption  of  oxygen  and 
conversion  into  ferric  hydrate  Fe2(HO)6. 

3.  (NH4)2CO3  or   Na2CO8  precipitates    white    FeCO3, 
which  rapidly  darkens  in  color. 

4.  K4Fe(CN)6  precipitates   white   K2Fe2(CN)6,    which 
rapidly  becomes  blue  by  oxidization  to  Fe5(CN)i2  (Prussian 
blue). 

5.  K3Fe(CN)6  precipitates  "Turnbull's  blue,"   Fe3Fe2 


6.  KCNS  produces  no  coloration. 

7.  BaCO3  produces  no  precipitate  in  cold  solution  of 
ferrous  salts. 

8.  Ferrous    compounds   are   converted  into   ferric  by 
oxidizing  agents,  such  as  HNO3,  KC1O3,  HC1,  Cl. 

9.  Fused  with  borax  in  the  oxidizing  flame,  yellowish- 
red  beads  are  produced  ;  in  the  reducing  flame  the  beads 
become  green. 

Ferric  Salts,  Feiv.     Solution,  Fe2Cl«. 

1.  H2S  in  acid  solutions  produces  a  precipitate  of  sul- 
phur, and  the  salt  is  reduced  to  protosalt  :  — 

Fe2Cle  +  H2S  =  2  FeCl2  +  2  HC1  +  S. 

2.  (NH4)2S   precipitates  black  FeS  and  S,  soluble  in 
HC1  and  HNO3. 

3.  (NH4)HO     or     KHO     precipitates     reddish-brown 
Fe2(HO)e,  insoluble  in  excess  of  reagent. 

4.  K4Fe(CN)6  precipitates  "  Prussian  blue,"  Fe«(CN)u, 
insoluble  in  HC1,  soluble  in  C2H2O4. 

5.  K3Fe(CN)6  produces  a  reddish-brown  color. 


REACTIONS  OF  THE  METALS  OF  THE  IRON  GROUP.       57 

6.  KCNS  produces,  even  in  dilute  solutions,  an  intense 
blood-red   color,   forming   a   soluble   iron   sulphocyanidc. 
HC1  does  not  destroy  coloration,  but  it  is  destroyed  by 
C2H3O2Na,  HgCl2,  H3PO4,  and  C4H6O6. 

7.  BaCO3  precipitates   ferric    solutions   completely   as 
Fe2(HO)6  mixed  with  basic  salt. 

8.  The  blow-pipe  reactions  are  the  same  as  for  ferrous 
compounds. 

9.  Heated  on  charcoal  with  Na2CO3,  compounds  of  Fe 
yield  magnetic  particles,  but  no  globules. 

Characteristic  reactions :  ferrous  salts,  5 ;  flrric,  4,  6. 

67.   Nickel,  Ni"' iv,  58.7.     Solution,  NiN2O6. 

1.  (NH4)2S  precipitates  black  NiS,  slightly  soluble  in 
excess,  forming  a  brown  solution,  from  which  NiS  is  pre- 
cipitated on  boiling,  insoluble  in  HC1,  soluble  in  HNO3 
and  aqua  regia. 

2.  (Na)HO  or  KHO  precipitates  light-green  Ni(HO)2, 
insoluble  in  excess  of  the  reagent. 

3.  (NH4)HO  produces  a  precipitate  of  Ni(HO)2,  readily 
soluble  in  excess,  yielding  a  blue  fluid.     Acid  solutions, 
or  those  containing  salts  of  ammonia,  yield  no  precipitate 
with  (NH4)HO. 

4.  KCN  precipitates  yellowish-green  Ni(CN)2,  soluble 
in  excess  and  reprecipitated  by  HC1  or   H2SO4,   and  if 
boiled  with  a  strong  solution  of  NaCIO,  yields  a  black  pre- 
cipitate of  Ni2(HO)6. 

5.  KNO3  in  presence  of  C2H4O2  produces  no  precipitate. 

6.  Fused  with  borax  compounds  of  Ni,  yields  reddish- 
yellow  beads  when  hot  in  oxidizing  flame.     In  reducing 
flame  the  bead  becomes  gray. 

Characteristic  reactions,  2,  3. 


58  ELEMENTARY  QUALITATIVE  ANALYSIS. 

68.    Cobalt,  Co"' iv,  58.7.     Solution,  CoN2O6. 

1.  (NH4)2S  precipitates  black  CoS,  insoluble  in  excess 
of  reagent  and  in  HC1,  soluble  in  aqua  regia. 

2.  KHO  or  NaHO  precipitates  blue  basic  salts,  which 
turn  green  on  exposure  to  air  by  oxidation. 

3.  (NH4)HO  precipitates  the  same  as  above,  soluble  in 
excess,  yielding  a  reddish-brown  fluid;  reprecipitated  by 
NaHO  or  KHO.     Acid  solutions,  or  those  containing  salts 
of  ammonia,  are  not  precipitated. 

4.  KCN   precipitates  light-brown  Co(CN)2,  soluble  in 
excess  of  the  reagent  by  formation  of  2KCN,--Co(CN)2. 
This  solution   is  reprecipitated   by  addition   of   HC1   or 
H2SO4.     If  to  the  solution  in  excess  of  KCN  a  few  drops 
of  HC1  be  added,  and  the  solution  boiled  for  some  time, 
K3Co(CN)6,  potassium  cobaltic  cyanide,  is  formed,  which  is 
not  reprecipitated  by  HC1  or  H2SO4,  nor  by  NaClO. 

5.  KNO2   added  to  cobalt  solutions    with   addition   of 
acetic  acid  precipitate,  on  standing,  a  yellow  crystalline 
double  salt. 

6.  Compounds  of  cobalt  fused   with   borax   in   either 
blow-pipe  flame  yield  deep-blue  beads. 

69.   Zinc,  Zn",  65.2.     Solution,  ZnSO4. 

1.  (NH4)2S  precipitates  white  ZnS,  insoluble  in  excess 
of  reagent  and  in  KHO ;  soluble  in  the  mineral  acids. 

2.  KHO  or  NaHO  precipitates  white  Zn(HO)2,  soluble 
in  excess  of  either  reagent  and  in  (NH4)HO ;  reprecipi- 
tated from  dilute  solutions  by  boiling,  but  not  by  NH4C1. 

3.  Na2C03  precipitates  white  basic  carbonate,  insoluble 
in  excess  oi  reagent. 

4.  (NH4)2CO3  precipitates  also  the  basic  carbonate,  but 
it  is  soluble  in  excess  of  the  reagent. 


REACTIONS  OF  THE  METALS  OF  THE  IKON  GROUP.      59 

5.  Heated  on  charcoal  with  Na2COa   in    the   reducing 
blow-pipe  flame,  a  yellow  incrustation  of  ZnO  is  obtained, 
which  becomes  white  when  cold. 

6.  Heated   on   charcoal   by  the   blow-pipe  flame,  after 
moistening  with  CoCl2  solution,  an  infusible  green  mass  is 
obtained. 

Characteristic  reaction,  1. 

7O.    Aluminium,  Aliv,  27.3.     Solution,  A12.3SO4. 

1.  (NH4)2S  precipitates  white  A12(HO)6,  soluble  in  acids. 

2.  KHO    or    NaHO    produces   also    a    precipitate    of 
A12(HO)6,  soluble  in  acids,  even  in  hot  acetic  acid,  and  in 
excess  of  the  reagent.     This  solution  is  not  precipitated 
by   H2S,    but  is  reprecipitated   by  NH4C1,  or   by   adding 
(NH4)HO  after  acidifying  with  HCL 

3.  (NH4)IIO  also  precipitates  A12(HO)6,  soluble  in  a 
very  large  excess  of  the  reagent,  more  difficultly  soluble 
in  presence  of  salts  of  ammonia. 

4.  BaC03  produces  a  precipitate  of  A12(HO)6  mixed 
with  basic  salt. 

5.  Na2HP04  precipitates  aluminium  phosphate,  insolu- 
ble in  (NH4)HO  and  in  NH4C1,  but  soluble  in  KHO  or 
NaHO,  and  in  acids.     It  does  not,  however,  dissolve  in 
hot  acetic  acid  like  aluminium  hydrate. 

6.  Heated   on    charcoal  in  the  blow-pipe    flame,   then 
moistened  with  CoCl2,  and  reheated,   an   infusible    blue 
mass  is  obtained. 

Characteristic  reactions,  2,  6. 

71,   Manganese,  Mn"' iv,  55.     Solution,  MnSO4. 

1.    (NH4)2S  produces  a  flesh-colored  precipitate  of  MnS, 
soluble  in  acids,  even  in  acetic  acid. 


60  ELEMENTARY  QUALITATIVE  ANALYSIS. 

2.  KHO  or  NaHO  produces  a  dirty-white  precipitate  of 
Mn(HO)2,  insoluble  in  excess  of  the  reagent;  the  pre- 
cipitate rapidly  darkens  in  color  by  absorption  of  oxygen. 
The   freshly-precipitated  hydrate  is  dissolved  by  NH4C1, 
but  the  higher  oxide  is  insoluble. 

3.  (NH4)HO     produces     the     same     precipitate     of 
Mn(HO)2,  insoluble  in  excess  of  the  reagent;  but  it  gives 
no  precipitate  if  the  manganese  solution  contain  NH4C1, 
Such  a  solution  on  standing  precipitates  the  dark-brown 
hydrate. 

4.  Na2CO3  produces   a   white    precipitate   of    MnCOs, 
which  darkens  in  color  by  absorption  of  oxygen. 

5.  If  any  manganese  solution  (free  from  chlorine)  be 
treated  with  PbO2  and  then  boiled  with  HNO3,  it  is  con- 
verted into  permanganate,  which  is  recognized  by  its  pink 
color  as  soon  as  the  mixture  has  settled. 

6.  If  any  manganese  compound  be  fused  on  platinum 
foil  with  Na2CO3  and  a  trace  of  KNO3,  it  is  converted  into 
Na2MnO4,  recognized  by  its  bright-green  color. 

7.  Fused  with  borax  in  the  oxidizing  flame,  an  amethyst- 
colored  bead  is  obtained,  which  becomes  colorless  in  the 
reducing  flame. 

Characteristic  reactions,  1,  6,  7. 

72.    Chromium,  Criy,  52.1.     Solution,  Cr2.3SO4. 

1.  (NH4)2S    produces    a    bluish -green    precipitate    of 
Cr2(HO)6,  insoluble  in  excess  of  the  reagent,  soluble  in 
acids. 

2.  (NH4)HO  also  precipitates  the  hydrate,  soluble  to 
some  extent  in  excess,  yielding  a  pink  fluid,  but  on  heat- 
ing, the  precipitation  is  complete. 

3.  KHO  or  NaHO  precipitates  also  Cr2(HO)6,  soluble 


REACTIONS  OF  THE  METALS  OF  THE  IRON  GROUP.      61 

however  in  excess,  yielding  a  green  or  bluish-violet  solu- 
tion. On  continued  boiling  or  addition  of  NH4C1  and 
heating,  the  hydrate  is  reprecipitated. 

4.  BaCO3  produces   a   precipitate   of   Cr2(HO)6  along 
with  basic  salt ;  the  precipitation  is  not  complete  till  the 
mixture  has  stood  some  time. 

5.  Fused  with   Na2CO3  and   KNO3  on   platinum   foil, 
yellow  Na2CrO4  is  obtained. 

6.  Fused  with  borax  in  either  flame  (but  best  in  the 
reducing  flame),  green  beads  are  obtained. 

Characteristic  reactions,  color  of  solutions  and  bead. 


62 


ELEMENTARY   QUALITATIVE   ANALYSIS. 


TABLE 

73.  Iron  Group  (III.).  Separation  of  Iron,  Nickel, 
To  filtrate  from  the  sulphides  of  the  Cu  and  As  groups  add  (NHJHO  (till 
shake  for  some  time.  Filter.  Wash  well  with  H2O,  containing  (NH4)2S, 

dilute  HC1, 
RESIDUE. 

NiS  and  CoS. 

Test  for  Co  by  borax  bead.  Dissolve  the  black  residue  in  HC1  and 
KC1O3.  Boil  down  just  to  dryness,  dilute  with  H20,  add  KCN  in  excess, 
then  a  drop  of  acetic  acid,  boil  for  a  few  minutes,  add  NaCIO  in  excess, 
and  boil  again.  A  black  precipitate  indicates  Nickel.  The  filtrate  from 
this  precipitate  may  be  tested  for  Co  by  evaporating  to  dryness,  and  fusing 
in  a  borax  bead.  Blue  color  indicates  Cobalt. 

METHOD  I.     Cr  is  absent. 

Boil  down  with  a  little  KC103  till  it  smells  of  Cl. 
strongly  alkaline.    Filter. 


Add  pure  NaHO  till 


RESIDUE. 


FILTRATE. 


Fe.2(HO)6,  Mn(HO)2. 

Wash    with  hot  H20,   dissolve   in 

HC1,  add  (NHJHO,  and  filter. 

RESIDUE.        I       FILTRATE. 


A  white 
Confirm 


Fe2(HO)6. 

Mn. 

Dissolve       in 

Boil  down  and 

HC1.     Test  with 

ignite,   to  expel 

K4Fe(CN)c.Blue 

salts   of    ammo- 

precipitate indi- 

nium.      Fuse 

cates  Iron.  To  as- 

with NaHO  and 

certain   whether 

KN03.    A  green 

the  iron  is  pres- 

residue indicates 

ent  as  ferrous  or 

Manganese. 

ferric    salt,    the 

Traces  of  Ni  and 

original     solution 

Co     are    found 

must    be   tested 

along    with    the 

with  K4Fe(CN)6, 

Mn. 

and  K3Fe(CN)6. 

Al,  Zn. 
Divide  into  two  parts. 

1.  Add  H2S  or  (NHJ2S. 
precipitate  indicates  Zinc. 
by  flame  reaction. 

2.  Add     HC1     till     acid,     then 
(NHJHO   till  alkaline.      A  white 
precipitate    indicates    Aluminium. 
Confirm  by  flame  reaction 


KEACTIONS  OF  THE  METALS  OF  THE  IKON  GROUP.   63 


V. 

Cobalt,  Aluminium,  Zinc,  Manganese,  and  Chromium. 

alkaline)  +  NH4C1  +  (NH4)2S.  Warm  the  mixture  gently  in  a  small  flask  and 
and  finally  once  with  H20  alone.  Treat  the  precipitate  with  cold 
and  filter. 

FILTRATE. 


Or,  Al,  Fe,  Zn,  and  Mn. 

(Green  or  violet  if  Cr  be  present.     Boil  down  a  portion,  and  test  for  Cr  by 
borax  bead.     Adopt  Method  I.  if  absent,  Method  II.  if  present.) 


METHOD  II.     Cr  is  present. 

Boil  down  with  a  little  KC1O3  till  it  smells  of  Cl.  Add  Na2C03  or  NaHO 
till  just  neutral  or  slightly  acid;  allow  to  become  perfectly  cold.  Add  ex- 
cess of  BaC03,  place  in  a  flask,  cork  up  and  shake  well,  allow  to  stand  till 

clear.     Filter. 
RESIDUE.  FILTRATE. 


Fe2(HO)6,  Cr2(HO)6,  A12(HO)6  (also 

excess  of  BaC03). 
Wash  well,  boil  with  pure  NaHO, 
and  filter;  add  HC1  to  the  filtrate, 
and  then  (NHJHO  till  alkaline.    A 
white  precipitate  indicates  Alumin- 
ium.    Confirm    by  flame  reaction. 
Fuse  the  residue  insoluble  in  NaHO 
with  a  mixture  of  Na2C03  and  KNO3  ; 
extract  with  water,  and  filter. 
RESIDUE.  FILTRATE. 


1 
Fe2(HO)6. 
Dissolve       in 
HC1,    and    test 
withK4Fe(CN)6. 
A  blue   precipi- 
tate      indicates 
Iron. 

1, 

Yellow  in  col- 
or.  Acidify  with 
acetic  acid  ;  add 
lead  acetate.    A 
bright-yellow 
precipitate  indi- 
cates Chromium. 

Zn,  Mn. 

Precipitate  the  Ba  present  with 
H2S04  in  the  hot  solution.  Boil 
well,  and  filter;  add  NaHO.  Pre- 
cipitate indicates  Manganese.  Con- 
firm by  fusing  "with  Na2C03  and 
KN03  on  platinum  foil.  To  the 
filtrate  from  the  Mn(HO)2  add 
(NH4)aS.  A  white  precipitate  indi- 
cates Zinc.  Confirm  by  flame  re- 
action. 


64  ELEMENTARY   QUALITATIVE   ANALYSIS. 

Reactions  of  the  Metals  of  the  Barium  Group 

(Group  IV.). 

74.  Metals  whose  carbonates  are  insoluble  in  water, 
and  whose  solutions  are  precipitated  on  the  addition  of 
(NH4)2CO3.  As,  however,  the  carbonates  are  soluble  in 
acids,  the  solution,  if  acid,  must  be  neutralized  by  addi- 
tion of  (NH4)HO. 

BARIUM,    STRONTIUM,   CALCIUM. 

Barium,  Ba",  137.     Solution,  BaCl2. 

1.  (NH4)2CO3   precipitates    white    BaCO3,    soluble    in 
acids,  and  to  a  slight  extent  in  NH4C1. 

2.  K2CO3  or  Na2CO3  precipitates  also  BaCO3,  insoluble 
in  excess  of  either  reagent. 

3.  H2SO4  or  any  soluble  sulphate,  even  in  dilute  solu- 
tions, precipitates  heavy  white  BaSO4,  insoluble  in  acids, 
alkalies,  or  salts  of  ammonium. 

4.  CaSO4    or    SrSO4    precipitates    immediately    white 
BaSO4. 

5.  H2SiF6  precipitates  white  BaSiF6. 

6.  C2(NH4)2O4  precipitates   white   C2BaO4,   soluble   in 
HC1  and  in  HNO3. 

7.  K2CrO4  precipitates    yellow    BaCrO4,    insoluble   in 
C2H4O2,  but  soluble  in  HC1  and  HNO3. 

8.  Heated  in  the  lamp  flame,  a  green  coloration  is  pro- 
duced, especially  on  moistening  the  salt  with  HC1. 

Characteristic  reactions,  3,  8. 

75.   Strontium,  Sr",  87.5.     Solution,  SrCl2. 

1.  (NH4)2CO8  or  K2CO3  precipitates  white  SrCO3, 
soluble  in  acids,  but  less  soluble  in  NH4C1  than  BaCO3. 


REACTIONS  OP  THE  METALS  OF  THE  BARIUM  GROUP.     65 

2.  H^SO4  precipitates  white  SrSO4,  much  less  insoluble 
in  H2O  than  BaSO4;  it  therefore  precipitates  from  dilute 
solutions  only  on  standing  or  warming.     SrSO4  is  slightly 
soluble  in  HC1. 

3.  CaSO4,  after  standing  some  time,  precipitates  white 
SrS04. 

4.  H2SiF6  does  not  precipitate  strontium  solutions. 

5.  C2(NH4)2O4  precipitates   white    C2SrO4,    soluble   in 
HC1  and  in  HNO3,  also  to  a  slight  extent  in  NH4C1,  but 
very  sparingly  in  C2H4O2. 

6.  K2CrO4,  only  in  concentrated  solutions,  precipitates 
yellow  SrCrO4,  soluble  in  C2H4O2. 

7.  Heated  in  the  lamp  flame,  a  crimson  coloration  is 
produced,  especially  on  moistening  the  salt  with  HC1. 

Characteristic  reactions,  3,  7. 

76.    Calcium,  Ca",  40.     Solution,  CaCl2. 

1.  (NH4)2CO3  or  K2CO3  precipitates  white  CaCO3,  which 
becomes  crystalline  on  heating. 

2.  H2SO4  precipitates  from  strong  solutions  of  calcium 
salts  CaSO4  as  a  white  precipitate,  which  dissolves  in  a 
large  excess  of  water,  and  also  in  acids. 

3.  CaSO4  produces  no  precipitate. 

4.  H2SiF6  produces  no  precipitate. 

5.  C2(NH4)2O4,    even   in   dilute   solutions,  precipitates 
white  C2CaO4,  soluble  in  HC1  or  HNO3,  but  insoluble  in 
C2H2O4  or  in  C2H4O2. 

6.  Heated  in  the  lamp  flame,  a  dull-red  coloration  is 
produced,  especially  on  moistening   the   salt   with   HC1. 
This  reaction  is  imperceptible  in  presence  of  Ba  or  Sr  salts. 

Characteristic  reactions,  5,  6. 


0? 
I  J7EESIT7 


66  ELEMENTARY  QUALITATIVE  ANALYSIS. 


TABLE   VI. 

77.    Barium  Group  (IV.).     Separation  of  Barium,  Strontium,  and 

Calcium. 

Heat  filtrate  from  iron  group,  add  to  the  hot  solution  NH4C1  and 
(NH4)2C03,  and  filter.  Wash  precipitate  with  hot  H20,  dissolve  in  HC1, 
and  add  CaS04  solution.  An  immediate  precipitate  indicates  Barium; 
a  precipitate  after  some  time  indicates  Strontium,  or  a  dilute  Barium 
solution.  (Test  another  portion  with  SrS04for  Ba.)  To  another  portion  of 
the  solution  in  HC1  add  H2S04,  and  boil  to  remove  Ba  and  Sr.  Filter. 
Neutralize  filtrate  with  (NHJHO,  and  add  C2(NHJ204.  An  immediate 
precipitate  indicates  Calcium. 


TO  TEST  FOR   STRONTIUM. 

I.   Ba  present ;  Ca  absent. 

Dissolve   the  carbonate  in  HC1,  and  evaporate  to  dryness.    Treat  the 

residue  with  strong  alcohol.    Filter. 
RESIDUE.  FILTRATE. 

BaCL,.  Sr. 

Confirm  by   flame   test.     Green          Confirm  by  lighting  the  alcoholic 
Coloration  indicates  Barium.  solution.     Crimson  coloration  indi- 

cates Strontium. 


II.  Ba  absent ;  Ca  present. 

Dissolve  as  before,  and  precipitate  with  H2S04.    Filter,  and  wash  well. 
RESIDUE.  FILTRATE. 

r  n 

SrSO4.  Ca. 

If  small,  burn  the  filter  in  the  re-  Neutralize     the     solution     with 

ducing  gas  flame  to  convert  SrS04  (NHJ HO,  and  test  with  C2(NH4)20.. 

into  SrS;    moisten  with  HC1,  and  White    precipitate    indicates    Cal- 

test  in  the  lamp  flame.     Crimson  cium. 
coloration  indicates  Strontium. 


KE ACTIONS  OF  THE  METALS  OF  THE  IIARIUM  GliOCJF.     07 


ANOTHER  METHOD. 

Dissolve  in  HN03,  and  evaporate  to  dryness.     Treat  with  strong  alcohol. 

Filter. 
RESIDUE.  FILTRATE. 


Sr(NO3)2.  Ca. 

Confirm  as  above.  Confirm  as  above. 


III.  Ba  and  Ca  present. 

Add    H2S04  to   the   HC1   solution   (diluted  to   prevent  precipitation   of 

Calcium),  and  filter. 
RESIDUE.  FILTRATE. 


BaS04,  SrSO4.  Ca. 

Boil    in   a  beaker  with   a   little  Neutralize     the     solution     with 

water,   together  with  a  mixture  of       ( NH4)  HO,  and  test  with  C2(NHJ204. 
three    parts    K2SO4  and  one   part      White    precipitate    indicates    Cal- 
K2C03.     Filter,  and   treat  residue      cium. 
with  HN03.    The  SrS04  is  dissolved, 
and    the   BaS04    left    undissolved. 
(Traces  of  Ca  may  be  found  with 
the  ST.) 


ANOTHER  METHOD. 

Dissolve  the  carbonates  in  C2H402,  and  precipitate  the  Ba  with  K2Cr04 
Filter.  Precipitate  the  Sr  and  Ca  by  (NH4)2C03,  and  proceed  as  in 
Method  II.  (Ba  absent,  Ca  present). 


68  ELEMENTARY   QUALITATIVE   ANALYSIS. 

Reactions  of  the  Metals  of  the  Potassium  Group 

(Group  V.). 

78.  Metals  whose  solutions  are  unprecipitated  by  the 
preceding  group  reagents,  and  having  no  common  pre- 
cipitant. They  are  therefore  recognized  by  individual 
tests. 

MAGNESIUM,   POTASSIUM,   AMMONIUM,    SODIUM. 

Magnesium,  Mg",  24.     Solution,  MgSO4. 

1.  (NH4)HO  and  (NH4)2CO3  give  no  precipitates  in  the 
presence  of  salts  of  ammonium. 

2.  Na2HPO4  in  presence  of  (NH4)HO  and  (NH4)C1  a 
crystalline  white  precipitate  of  MgNH4PO4.     The  precipi- 
tation is  slow  from  dilute  solutions,  but  may  be  hastened 
by  stirring  with  a  glass  rod  and  warming.     Precipitate 
soluble  in  dilute  mineral  acids  and  in  C2H4O2,  almost  in- 
soluble in  dilute  solution  of  (NH4)HO. 

3.  H2SO4,   H2SiF6,    and    C2(NH4)2O4  give   no   precipi- 
tates. 

4.  Heated  on  charcoal  in  the  blow-pipe  flame,  and  then 
moistened  with  CoCl2  or  CoN2O6  and  reheated,  gives  a 
pink  mass. 

Characteristic  reactions,  2,  4. 

79.   Potassium,  K',  39.1.     Solution,  KC1. 

1.  PtCl4,  except  in  dilute  solutions,  gives  a  crystalline 
yellow  precipitate  of  2  KC1  -f  PtCl4.     The  precipitation  is0 
hastened  by  stirring  or  the  addition  of  alcohol. 

2.  H2Tr  or  NaHTr  precipitates  white  crystalline  KHTr 
from  concentrated  solutions. 

3.  H2SiF6  precipitates  white  gelatinous  K2SiF6. 


REACTIONS  OF  THE  METALS  OF  THE  POTASSIUM  GROUP.  69 

4.  Heated  on  platinum  wire,  potassium  compounds 
color  the  flame  violet,  appearing  reddish-violet  through 
blue  glass. 

Characteristic  reaction,  4. 

8O.   Ammonium,  NH4,  18.     Solution,  NH4C1. 

1.  PtCl4    gives    a    crystalline    yellow    precipitate    of 
2NH4C1  -f-PtC!4,  except  in  dilute   solutions.     Precipitate 
insoluble  in  alcohol  and  ether.     On  ignition,  precipitate 
leaves  a  residue  of  spongy  platinum. 

2.  NaHTr  or   H2Tr  give  in   strong  solutions  a  white 
precipitate  of  (NH4)HTr. 

3.  H2SiF6  gives  no  precipitate. 

4.  Nessler's   solution  gives  a  brown  precipitate,  or  in 
any  dilute  solutions  a  yellow  coloration. 

5.  Heated  with  NaHO  or  KHO,  compounds  of  ammo- 
nium evolve  ammonia  gas,  recognized  by  its  odor,  alkaline 
reaction,  and  fuming  with  HC1. 

6.  Heated  on  platinum  foil,  all  compounds  of  ammonia 
volatilize  completely. 

Characteristic  reactions,  4,  5,  6. 

81.    Sodium,  Na',  23.     Solution,  NaCl. 

1.  PtCl4,  NaHTr,  and  H2Tr  give  no  precipitates. 

2.  H2SiF6     gives     precipitate     of     white     gelatinous 
Na2SiF6. 

3.  The  salts  of  sodium  being  almost  without  an  excep- 
tion soluble  in  water,  the  flame  test  alone  serves  to  dis- 
tinguish the  metal.     Heated  on  platinum  foil  or  wire  in 
non-luminous  flame,  an  intense  yellow  color  is  produced, 
not  seen  when  viewed  through  blue  glass. 


70 


ELEMENTARY   QUALITATIVE   ANALYSIS. 


TABLE  VII. 

82.    Group  V.     Separation  of  Magnesium,  Potassium,  Sodium,  and 

Ammonium . 

The  filtrate  from  the  Barium  group  is  concentrated  by  evaporation, 
and  a  portion  ignited  on  platinum  foil.  If  no  residue  is  left  on  ignition, 
Mg,  K,  and  Na  are  absent. 


Detection  of           Detection  of 

Detection  of  K  and  Na. 

NH4.                         Mg. 

(1.)  Mg  being 

(2.)  Mg  being 

The     original         To   a   portion 

absent. 

present. 

substance  or  so-     oftheeoncentrat- 

Evaporate  an- 

Evaporate the 

lution  is  heated     ed  cold  solution 

other  portion  of 

solution   to  dry- 

with  NaHO  in  a     add      (NPLJHO 

the    solution    to 

ness,  ignite  resi- 

test-tube.    Prejs-     and      Na2HPO4. 

dryness,      ignite 

due,  dissolve  in 

ence  of  Ammo-     White     crystal- 

residue,  dissolve 

water,   and    add 

nium  shown  by     line    precipitate 

in  a  small  quan- 

baryta water  un- 

smell,    by     the      denotes  Magne- 

tity  of  water,  fil- 

til   the   solution 

white  fumes  with     slum. 

ter  if    required, 

has   an  alkaline 

HC1,  and  by  its 

and  add  to   the 

reaction;  boil;  fil- 

action on  red  lit- 

clear   liquid 

ter.    To  filtrate, 

mus  paper. 

PtCl4,  evaporate 

add    (NHJ2C03, 

nearly    to    dry- 

heat,  filter,  eva- 

To detect  Na. 

ness,     and     add 

porate    to     dry- 

Evaporate      alcoholic      solution 

alcohol.    Yellow 

ness,     and     test 

(which  must  have  a  yellow   color, 

precipitate  indi- 

the residue    for 

showing  that   excess  of  PtCl4  has 

cates  Potassium. 

K  and  Na. 

been  added)  nearly  to  dryness,  add 

a  grain  or  two  of  sugar,  and  ignite 

residue.    Exhaust  with  water,  fil- 

ter, evaporate  to  dryness  ;  and  if  a 

residue  be  left,  test  it  by  flame  re- 

action  for  Na.     Yellow  coloration 

indicates  Sodium. 

PRECIPITATIONS   IN   THE   FIVE   GROUPS   OF   BASES.     71 


Iodides  . 

Bromides 

Sulphites  , 

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Oxalates 

Phosphate 

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72  ELEMENTARY  QUALITATIVE  ANALYSIS. 

REACTIONS  OF  THE  ACIDS. 
s  • 

83.  Grouping  of   the  Acids.     The   acids   can  be  ap- 
proximately classified  by  means  of  certain  group  reagents. 

They  are  divided  into  two  great  classes :  inorganic  and 
organic  acids.  These  are  easily  distinguished  by  the 
action  of  heat. 

Salts  of  inorganic  acids,  when  heated  to  redness,  are  not 
charred ;  salts  of  organic  acids  are  at  once  charred,  owing 
to  decomposition  and  separation  of  carbon  (with  the  ex- 
ception of  acetic  and  formic  acids). 

84.  Grouping  of  the  Inorganic  Acids. 

Group  I.     {Sulphuric  Acid  Group.) 

Group  reagent,  BaCl2  in  presence  of  HC1. 
Sulphuric  acid,  hydrofluo-silicic  acid. 
The  acids  of  this  group  are  precipitated  by  BaCl2,  and 
the  precipitate  is  not  dissolved  on  addition  of  HC1. 

Group  II.     {Phosphoric  Acid  Group.) 

Group  reagent,  BaCl2. 

Phosphoric,  boric,  hydrofluoric,  carbonic,  silicic,  sul- 
phurous, arsenious,  arsenic,  iodic,  chromic  acids. 

The  acids  of  this  group  are  precipitated  in  neutral 
solutions  by  BaCl2. 

Group  III.     {Hydrochloric  Acid  Group.) 

Group  reagent,  AgNO3. 

Hydrochloric,  hydrobromic,  hydroiodic,  hydrocyanic, 
and  hydro  sulphuric  acids. 

The  acids  of  this  group  are  precipitated  by  AgNO3,  and 
not  by  BaCl2. 


REACTIONS   OF   THE   ACIDS.  73 

Group  IV.     (Nitric  Acid  Group.) 

Nitric,  chloric,  and  perchloric  acids. 
These  acids  are  not  precipitated  by  any  reagent,  as  all 
their  salts  are  soluble  in  water. 

Reactions  of  the  Inorganic  Acids  belonging  to 
Group  I. 

85.   Acids  precipitated  by  BaCl2  in  presence  of  HC1. 

SULPHURIC  ACID,   HYDROFLUO-SILICIC   ACID. 

Sulphuric  Acid,  H2SO4,  98. 

1.  BaCl2  precipitates  a  white  BaSO4,  insoluble  in  HC1 
or  HNO3.      In  very  dilute  solutions  the  precipitation  is 
not   immediate,  but   on   standing,   the  solution  becomes 
clouded,  and  ultimately  the  precipitate  subsides. 

2.  Pb(NO3)2  precipitates  a  heavy  white  PbSO4,  soluble 
in  NaHO,  and  in  boiling  HC1  (on  allowing  this  solution 
to  cool,  PbCl2  crystallizes  out). 

3.  fused  on    charcoal   with   Na2COs  in   the   reducing 
flame  of  the  blow-pipe,  a  sulphide  is  produced.     If  the 
fused  mass  be  moistened  with  HC1,  the  odor  of  H2S  is  at 
once  perceptible ;  or  if  it  be  placed  on  a  bright  piece  of 
silver  and  moistened  with  water,  a  black  stain  of  Ag2S  is 
produced. 

86.   Hydrofluo-silicic  Acid,  H2SiF6,  144. 

1.  BaCl2  precipitates  a  crystalline  BaSiF6,  insoluble  in 
HCL 

2.  KC1  precipitates  a  gelatinous  K2SiF6. 

3.  Heated  with  H2SO4  in  a  leaden  cri  cible  covered  with 


74  ELEMENTARY   QUALITATIVE   ANALYSIS. 

a  piece  of  glass,  the  latter  will  be  etched  by  the  evolved 
HF. 

.Reactions  of  the  Acids  belonging  to  Group  II, 

.    87.   Acids  precipitated  by  BaCl2  in  neutral  solutions. 

PHOSPHORIC,  BOEIC,  HYDROFLUORIC,  CARBONIC,  SILICIC, 
SULPHUROUS,  ARSENIOUS,  ARSENIC,  IODIC,  AND 
CHROMIC  ACIDS. 


Phosphoric  Acid,  HsPO^,  98. 

1.  BaCl2  precipitates  a  white  BaHPO4,  readily  soluble 
in  HNO3  or  HC1,  but  "with  difficulty  in  NH4C1. 

2.  Mg2SO4,  along  with  (NH4)HO  and  NH4C1,  precipi- 
tates a  white  crystalline  Mg(NH4)PO4  +  6  H2O,  insoluble 
in  (NH4)HO,  but  soluble  in  HC1,  HNO3,  and  acetic  acid. 
In  dilute  solutions  the  precipitation  does  not  take  place 
till  after  the  lapse  of  some  time,  but  is  promoted  by  stirring 
and  gentle  warming. 

3.  AgNO3  precipitates    a    yellow   Ag8PO4,   soluble   in 
HNO3,  and  also  in  (NH4)HO. 

4.  Lead  acetate  precipitates  a  white  Pb8(PO4)2,  soluble 
in  HNO3,  but  almost  insoluble  in  acetic  acid. 

5.  Fe2Cl6,  in  presence  of  excess  of  sodium  acetate,  pre- 
cipitates a  yellowish  FePO4,  soluble  in  HC1,  and  in  excess 
of  Fe2Cl6,  which  must.be  added  drop  by  drop. 

6.  Ammonium   molybdate  produces  in  solutions  acidi- 
fied by  HNO3  a  yellow  color,  and  then  a  precipitate  ;  this 
reaction  is  hastened  by  warming. 

88.   Boric  Acid,  B(HO)8,  62. 

1.   BaCl2  precipitates    a   white    Pa(BO2)2,   soluble    in 
acids. 


REACTIONS   OF   THE    ACIDS.  75 

2.  AgNOg  produces  in  strong  solutions  a  yellowish-white 
precipitate.     In  dilute  solutions  Ag2O  is  precipitated. 

3.  H2SO4  or  HC1,  added  to  hot  concentrated  solutions 
of  alkaline  borates,  on  cooling,  precipitates  a  crystalline 
B(HO),. 

4.  If  alcohol  containing  free  boric  acid  be  kindled,  it 
burns  with  a  green  flame,  best  seen  on  stirring  the  mix- 
ture.     Borates   may   be   examined   in   this  way  by  first 
adding  strong  H2SO4,  to  liberate  the  B(HO)3. 

5.  If  the  solution  of  a  borate  be  made  distinctly  acid 
with  HC1,  and  turmeric  paper  dipped  into  it,  the  latter,  on 
gentle  warming,  acquires  a  brown  tint,  which  is  turned 
blue  by  caustic  soda. 

89.   Hydrofluoric  Acid,  HF,  20. 

1.  BaCl2  precipitates  a  white  BaF2,  soluble  in  HC1,  and 
sparingly  in  NH4C1. 

2.  Ca012  produces  a  gelatinous  and  almost  transparent 
precipitate  of  CaF2,  made  more  apparent  on  addition  of 
(NH4)HO.     The  precipitate  is  very  difficultly  soluble  in 
II Cl,  even   on  boiling,  and  is  nearly  insoluble  in  acetic 
acid. 

3.  Heated  with  H2SO4,   all   fluorides   are   decomposed 
with  evolution  of  HF,  which  is  recognized  by  its  power  of 
etching  glass. 

4.  Heated  with  a  mixture  of  borax  and  HKSO4,  on  a 
loop  of  platinum  wire  in  the  non-luminous  flame,  BF3  is 
produced,  coloring  the  flame  green. 

90.    Carbonic  Acid,  H2CO3,  H20  4  CO2. 

1.   BaCl2    in    neutral    solutions    precipitates    a    white 
BaCO3,  soluble  in  acids  with  effervescence. 


76  ELEMENTARY  QUALITATIVE   ANALYSIS. 

2i  Treated  with  dilute  HC1,  all  carbonates  at  once 
evolve  COz  with  effervescence,  which  turns  lime  water  a 
milky  white  from  the  formation  of  CaCO3. 

91.   Silicic  Acid,  Si(HO)4,  96. 

1.  BaCl2  precipitates  a  white  SiBa2O4,  which  is  decom- 
posed on  addition  of  HC1,  and  Si(HO)4  separates  out  as  a 
gelatinous  precipitate. 

2.  HC1,  added  drop  by  drop  to  a  strong  solution  of  a 
silicate,  precipitates  a  gelatinous  Si(HO)4;  but  if  added 
to  a  dilute  solution  or  in  large  excess,  no  precipitate  is 
obtained  until  the  mixture  has  been  evaporated  to  dryness 
and  ignited,  when  SiO2  separates  out,  and  this  is  not  re- 
dissolved  on  addition  of  HC1. 

3.  Fused  with  Na2CO3  in  a  loop  of  platinum  wire  in  the 
non-luminous  gas-flame,  effervescence  occurs  from  the  dis- 
engagement of  CO2,  and  the  bead  is  transparent  on  cool- 
ing, unless  the  Na2CO3  be  in  excess. 

4.  Fused  with  microcosmic  salt  on  a  loop  of  platinum 
wire  in  the  non-luminous  gas-flame,  solution  does  not  take 
place,  but  the  silica  floats  about  on  the  bead  undissolved. 

92.   Sulphurous  Acid,  H2SO3,  82. 

1.  BaCl2  precipitates  a  white   BaSO3,  soluble  in  HC1. 
On  addition  of  chlorine  water,  gives  a  white  precipitate  of 
BaSO4,  the  sulphite  being  oxidized  to  the  sulphate. 

2.  AgNO3  gives  a  white  precipitate  of  AgSO3,  dark- 
ened on  heating. 

3.  Added  to  a  mixture  of  Zn  and  HC1,  H2S  is  pro- 
duced, and  recognized  by  its  smell  and  by  its  action  on 
paper  moistened  with  a  solution  of  a  lead  salt,  blacken- 
ing it. 


REACTIONS   OF   THE  ACIDS.  .  77 

4.  H2S    decomposes   pure    H2SO3,    with   separation   of 
sulphur. 

5.  H2SO3  is  decomposed  by  HC1,  with  evolution  of  SO2. 


Y 


93.   Arsenious  Acid,  H3AsO3,  126. 

1.  AgNO3  gives  in  neutral  solutions  a  yellow  precipi- 
tate of  Ag3AsO3,  soluble  in  (NH4)HO. 

2.  MgSO4  +  (NH4)C1  +  (NH4)HO  give  no  precipitate.  • 
*A     3.   H2S  precipitates  As2S3,  yellow. 

94.   Arsenic  Acid,  H3AsO4,  142. 

1.  AgNO3  gives  in  neutral  solutions  a  light-brown  pre- 
cipitate of  AgAsO4. 

2.  MgSO4+(NH4)Cl  +  (NH4)HO    give    a  white   pre- 
cipitate of  MgNH4AsO4. 

3.  H2S  precipitates  AsgSa,  yellow. 


95.  lodic  Acid,  HIO8,  176. 

1.  BaCl2  gives  a  white  precipitate  of  BaI2O6,  soluble  in 
HN03. 

2.  AgNO3  precipitates  white  crystalline  AgIO3,  easily 
soluble  in  (NH4)HO,  but  sparingly  so  in  HNO3. 

3.  SO2  gives  at  first  a  precipitate  of  I,  which  is  con- 
verted into  HI  on  addition  of  excess  of  reagent. 

X  4.   HIO3  is  decomposed  by  H2S,  with  formation  of  an 
iodide  and  separation  of  S. 

5.   lodate  salts,  on  heating,  are  decomposed,  oxygen  being 
evolved.   In  some  cases  iodine  is  given  off  in  violet  vapors. 

96.    Chromic  Acid,  H2CrO4,  118.2. 

1.   BaCl2  precipitates  a  yellow  BaCrO4,  soluble  in  HC1 
and  HNO3,  but  insoluble  in  acetic  acid. 


78  ELEMENT AKY   QUALITATIVE   ANALYSIS. 

2.  H2S  in   presence   of  HC1   reduces  the   solution   to 
Cr2Cl6  (green),  with  separation  of  S.    In  neutral  solutions, 
Cr2(HO)  is  precipitated  along  with  S. 

3.  SO2  reduces  solutions  of  chromates  to  the  chromic 
salt,  the  color  of  which  is  green.     Chromates  are  likewise 
reduced  by  zinc  and  a  dilute  acid,  by  oxalic  acid  and 
dilute  sulphuric  acid,  by  strong  H2SO4,  by  strong  HC1, 

"and  by  boiling  the  solution,  acidified  with  HC1  or  H2SO4, 

along  with  alcohol. 
/x  4.   AgNO3  precipitates  a  dark-red  Ag2CrO4,  soluble  in 

HN03  and  in  (NH4)HO. 

5.  Lead  acetate  produces  a  bright-yellow  precipitate  of 
PbCrO4,  soluble  in  NaHO,  but  soluble  with  difficulty  in 
dilute  HNO3. 

6.  H2CrO4  is  precipitated  by  (NH4)2S  as  Cr2(HO)6. 

.Reactions  of  the  Acids  belonging  to-  Group  III. 

97.   Acids  precipitated  by  AftN^,  an^  "n*  ^y  ^nni- 

HYDROCHLORIC,    HYDROBROMIC,   HYDRIODIC,   HYDRO- 
CYANIC,   AND   HYDROSULPHURIC   ACIDS. 

Hydrochloric  Acid,  HC1,  36.5. 

1.  AgNO3  precipitates  a  white  curdy  AgCl,  which  be- 
comes violet  on  exposure  to  light.     The  precipitate  is  in- 
soluble in  HNO3,  but  soluble  in  (NH4)HO,  in  KCN,  in 
Na2S2O3,  and  also  to  some  extent  in  NaCl. 

2.  Heated    with    H2SO4    and    MnO2,    chlorides    yield 
chlorine  gas,  recognized  by  its  smell,  bleaching  action,  and 
green  color. 

3.  Dry  chlorides,  when  heated  in  a  retort  with  H2SO4 
and   K2Cr2O7,    yield    CrO2Cl2    (chromium    oxychloride), 


REACTIONS    OF    THE   ACIDS.  79 

which  distils  over  into  the  receiver  as  a  dark-red  liquid, 
decomposed  by  addition  of  water  or  (NH4)HO,  yielding  a 
yellow  solution,  which,  on  addition  of  a  lead  salt,  gives  a 
yellow  precipitate  of  PbCrO4. 

98.   Hydrolromic  Acid,  HBr,  81. 

1.  AgNO3  precipitates  a  pale-yellow  AgBr,  insoluble  in 
dilute  HNO3,  soluble  in  strong  (NH4)HO,  and  readily  in 
KCN  and  Na2S2O3. 

2.  TTpaf.prl  wit.h_  T-TpS^X  and  MnQ9,  bromides  yield  red 
vapors  of  Br,  recognized  by  its  powerful  odor. 

3.  Heated  in  a  retort  with  K2Cr2O7   and   H2SO4,   dry 
bromides  yield  dark-red  vapors,  which  condense  in  the 
receiver  to  a  liquid  of  the  same  color,  which  consists  of 
pure  bromine,  and   is   decolorized   on   adding   excess   of 
(NH4)HO. 

99.   Hydriodic  Acid,  HI,  128. 

AgNO3  precipitates  a  pale-yellow  Agl,  insoluble  in 
dilute  HNO3,  and  very  difficultly  soluble  in  (NH4)HO, 
but  readily  in  KCN  and  Na2S2O3. 

2.  Cuprous   sulphate   precipitates  a  dirty-white    Cu2I2, 
which  separates  most  completely  if  the  solution  be  made 
slightly  alkaline  with  Na2CO3. 

3.  KNO2  produces  no  reaction  in  solutions  of  iodides 
until  a  few  drops  of  HC1  or  H2SO4  are  added,  when  iodine 
is  at  once  liberated  and  colors  the  solution  yellow.     If  a 
little  starch  solution  be  now  added,  a  deep-blue  coloration 
results  from  the  formation  of  starch  iodide. 

4.  Chlorine   water  (or  the  gas)  liberates  iodine  from 
iodides,  but  excess  of  Cl   causes  the  formation  of  IC18, 


80  ELEMENTARY  QUALITATIVE   ANALYSIS. 

which  is  colorless,  and  gives  no  blue  coloration  with  starch 
solution. 

5.   Heated  with  MnO2  and  dilute  H2SO4,  violet  vapors  of 
iodine  are  obtained,  which  color  starch  paper  blue. 

1OO.    Hydrocyanic  Acid,  HCN,  27. 

*  1.  AgNO3  precipitates  a  white  AgCN,  insoluble  in 
HNO8,  with  difficulty  in  (NH4)HO,  but  readily  in  KCN 
and  Na2S2O8.  AgCN  is  decomposed  on  ignition,  and 
metallic  Ag  remains;  this  serves  to  distinguish  it  from 
AgCl,  which  is  not  decomposed  on  ignition. 

2.  If  a  solution  of  FeSO4  which  has  become  oxidized 
by  exposure  to  the  air,  be  added   to    the   solution   of  a 
cyanide   made  alkaline  with  NaHO,  a  bluish-green  pre- 
cipitate is  formed,  which  is  a  mixture  of  Prussian  blue 
with  the  hydrated  oxides  of  iron.     On  adding  HC1,  these 
last  are  dissolved,  and  the  blue  precipitate  remains. 

3.  HC1  decomposes  nearly  all  cyanides  with  evolution 
of  HCN,  recognized  by  its  odor,  resembling  bitter  almonds. 

4.  Hg(CN)2  cannot  be  detected  by  the  above  methods. 
The  dry  substance  is  detected  by  igniting  in  a  small  tube, 
when  cyanogen  gas  is  evolved,  or  the  solution  is  decom- 
posed by  H2S  and   filtered   from   the  HgS :    the   filtrate 
contains  HCN. 

1O1.    Hydrosulphuric  Acid  {Sulphuretted  Hydrogen}, 
H2S,  34. 

1.  AgNOg  precipitates  a  black  Ag2S,  insoluble  in  dilute 
acids. 

2.  Lead  acetate,  even  when  highly  dilute,  precipitates 
black  PbS. 


REACTIONS    OF   THE   ACIDS. 

3.  Sodium  nitro-pmsside,  in  presence  of  NaHO,  produces 
a  reddish-violet  coloration,  even  in  very  dilute  solutions. 
The  color  disappears  in  a  short  time. 

4.  HC1  or  H2SO4  decomposes  most  sulphides  with  evo- 
lution of  H2S,  recognized  by  its  disagreeable  odor  and  by 
its  blackening  paper  moistened  with  solution  of  lead. 

Reactions  of  the  Acids  of  Group    IV. 

1O2.   Acids  not  precipitated  by  any  reagent. 
Nitric  Acid,  HNO3,  63. 

1.  Nitrates  when  heated  evolve  oxygen,  and  in  some 
cases  nitrous  vapors  also.     On  fusing  a  nitrate  and  adding 
a  fragment  of  charcoal,  vivid  deflagration  occurs. 

2.  Free  HNO3  heated  with  Cu  gives  red  fumes ;  boiled 
with  pieces  of  silk  or  wool  turns  them  yellow. 

3.  If  to  a  solution  of  a  nitrate  FeSO4  and  concentrated 
H2SO4  be  poured  carefully  into  the  test-tube,  a  dark  ring 
will  appear  on  top  of  the  H2SO4,  which  will  be  violet,  red, 
or  dark-brown  according  to  the  quantity  of  HNO3  present. 
The  ring  disappears  on  warming. 

103.    Chloric  Acid,  HC1O3,  84.5. 

1.  H2SO4  decomposes  chlorates  with  evolution  of  C12O4, 
a  greenish-yellow  gas  having  a  powerful  odor.     If  heated, 
violent  explosions  occur ;  the  mixture  ought  therefore  to 
be  kept  cold,  and  only  very  small  quantities  should  be 
used. 

2.  When  chlorates  are  heated,  oxygen  is  evolved,  and  a 
metallic  chloride  remains,  which  may  be  dissolved  in  water, 
and  precipitated  as  AgCl  by  AgNO8. 


82  ELEMENTARY  QUALITATIVE   ANALYSIS. 

3.  Chlorates  are   reduced   by    SO2   with   liberation   of 
chlorine  or  its  oxides ;  hence  if  the  solution  of  a  chlorate 
be  colored  blue  with  indigo  solution,  it  is  decolorized  on 
adding  H2SO4  and  solution  of  Na2SO8.     (Distinction  from 
perchlorates.) 

4.  HC1  decomposes  chlorates  with  evolution  of  Cl  and 
C12O4,  a  mixture  called  euchlorine. 

5.  Heated  with  charcoal,  chlorates  deflagrate  violently. 

1O4.   Perchloric  Acid,  HC1O4,  100.5. 

1.  H2SO4  does  not  act  upon  perchlorates  in  the  cold, 
and  on  heating,  white  fumes  of  HC1O4  are  given  off,  but 
no  explosions  occur. 

2.  KC1  in  strong  solutions  precipitates  a  white  KC1O4. 

3.  Indigo  solution  is  not  decolorized  when  added  to 
perchlorates    warmed    with    HC1,   as    euchlorine   is   not 
evolved. 

4.  Dry  perchlorates  evolve  oxygen  on  heating. 

5.  Perchlorates  are  not  reduced  by  SO2. 


REACTIONS   OF   THE   ACIDS.  83 

TABLE   VIII. 

Detection  of  Inorganic  Acids  in  Mixtures. 

(The  following  acids  are  found  in  examination  for  bases,  which  ought 
always  to  precede  examination  for  acids  :  — 

H2SO3, 

H2S2O3, 

H2C<>3, 

H2S. 

Si(HO)4, 

H2Cr04, 

H3As03, 

H3As04.) 

I. 

Acids  in  Soluble  Bodies. 

1.  Neutralize  a  portion  of  the  solution  with  (NH4)HO,  and  add  BaCl2 
(or  Ba(NO3)2  if  Ag,  Hg2,  or  Pb  be  present)  :  precipitate  indicates  H2S04, 
H3PO4,  H3AsO3,  H3As04,  Si(HO)4,  H2CrO4,  and  large  quantities  of  B(HO)3 
and  HF.* 

To  precipitate,  add  H20,  and  then  HC1:  if  a  precipitate  remain,  H2S04 
was  present. 

2.  To  another  portion  of  the  neutralized  solution  add  AgN03 :  a  precip- 
itate indicates  one  or  more  of  these  acids ;  i.e., 

(a)  HC1,  HBr,  HI,  HCn,  H4Fe(Cn)6,  H3Fe(Cn)6,  H2S. 
(6)   H3PO4,  H3As04,  H3As03,  H2CrO4,  Si(HO)4,  B(HO)3.* 
To  the  precipitate  add  cold  dilute  HN03.     Acids  under  (a)  are  insol- 
uble; those  under  (6),  soluble. 

Detection  of  Acids  under  (a). 

To  a  portion  of  the  solution  add  starch  paste  and  one  drop  of  a  solution 
of  N./)3  in  H2SO4.  Blue  coloration  indicates  HI.  Add  chlorine  water  till 
the  blue  color  disappears,  and  shake  with  chloroform.  Reddish  brown 

*  Oxalic,  citric,  and  tartaric  acids  will  also  be  shown,  if  present. 


84  ELEMENTARY   QUALITIVE   ANALYSIS. 

color  indicates  the  presence  of  HBr.  HC1  is  detected  in  presence  of  the 
others  by  boiling  down  the  solution  to  dryness  and  distilling  residue  with 
K2Cr207  and  H2SO4  (see  97,  3). 

Detection  of  Acids  under  (6). 
Test  separately  for  each  acid  by  the  methods  already  given. 

Separation  of  H3As03,  H3As04,  and  H3P04. 

Acidify  solution  with  HC1,  add  Na2S03,  and  heat  until  no  smell  of  S02 
is  given  off. 

Pass  H2S  through  the  hot  solution,  filter,  and  test  for  H3P04  with  am- 
monium molybdate  :  yellow  precipitate  indicates  H3P04. 

Precipitate  another  portion  with  magnesia  mixture,  and  test  both  pre- 
cipitate and  filtrate  for  arsenic. 

Test  for  the  other  acids  by  the  following  reactions,  given  under  each 
acid :  — 

For  HCn,  by  test  3,  100. 

For  H2S,  by  test  4,  101. 

For  HN03,  by  tests  2  and  2,  102. 

For  HC1O3,  by  tests  1  and  2,  103. 

For  B(HO)3,  by  tests  4  and  5,  88. 

For  Si(HO)4,  by  tests  2  and  4,  96. 

For  H2S03,  by  test  3,  92,  and  smell  of  SO2  on  adding  HC1. 

For  CO2,  by  test  2,  90. 


APPENDIX. 


To  determine  the  empirical  formula  of  a  substance  from  its 
percentage  composition. 

RULE. 

Divide  the  percentage  amount  of  each  constituent  by  its  corre- 
sponding atomic  weight;  then  divide  each  quotient  so  found  by  the 
lowest  number}  and  reduce  them  to  their  simplest  ratios. 

EXAMPLES. 

1.  A  body  on  analysis  yielded  the  following  percentage 
composition:  — 

Carbon,    27.273  % 

Oxygen,  72.727 
100.000 

Calculate  its  formula. 
The  atomic  weight  of  carbon  is  12. 
The  atomic  weight  of  oxygen  is  16. 


Then,  C  =  =  2.2727  ; 

1^ 

0  =  1^21  =  4.5454. 
16 

Simplest  ratio  between  the  carbon  and  oxygen  is  as  1  :  2  ; 
for  2.2727  :  4.5454  :  :  1  :  2. 

Hence  the  formula  is  C02. 


94  APPENDIX. 

2.  A  compound  was  found  to  have  the  following  percentage 
composition  :  — 

Nitrogen,     82.353 

Hydrogen,    17.647 

100.000 

Calculate  its  formula. 

The  atomic  weight  of  nitrogen  =  14,  and  of  hydrogen  =  1. 

v      82.353 

JN  =  —      —  = 
14 

H  =  Hl= 


The  simplest  ratio  between  nitrogen  and  hydrogen  is  as  1  :  3  ; 
for  5.882  :  17.647  :  :  1  :  3. 

The  formula  of  the  body  therefore  is  NH3. 

3.  A  compound  of  iron  and  oxygen  has  the  following  per- 
centage composition  :  — 

Iron,          70.01 

Oxygen,    29.99 

100.00 

Calculate  its  formula. 

d 
Atomic  weight  of  iron  56.0,  and  of  oxygen  16.0. 

4.  Deduce  the  formulae  of  the  following  substances  :  — 

Nitrogen,    30.43 
Oxygen,       69.57 
100.00 


APPENDIX.  95 

5.  Potassium,  28.73 
Hydrogen,  0.73 
Sulphur,  23.52 
Oxygen,  47.02 

100.00 

6.  Carbon,  20.00 
Oxygen,          26.67 
Sulphur,         53.33 

100.00 

Relative  Density  of  Solids,  Liquids,  and  Gases; 
Vapor  Density. 

The  specific  gravity  (sp.  gr.),  or  relative  density  of  a  solid 
or  liquid  substance,  is  the  ratio  of  its  mass  to  the  mass  of  an 
equal  volume  of  some  liquid  taken  as  unity. 

The  standard  universally  adopted  is  pure  water  at  its  maxi- 
mum density.  The  number  expressing  the  relative  density  of 
a  solid  or  liquid  expresses  how  much  heavier  or  lighter  the 
substance  -is  than  an  equal  volume  of  water  at  4°  C. 

The  relative  density  of  a  solid  is  generally  ascertained  by 
the  following  formula,  the  body  being  first  weighed  in  air  and 
then  in  water  at  4°  C.,  and  the  weights  carefully  taken. 

Eel.  dens.  =  -       Weight  of  substance  (W) 

Weight  of  equal  vol.  of  water  at  4°  C. 

0 

W 


where  W  =  Weight  of  substance  in  water  at  4°  C. 

If  the  solid  be  lighter,  volume  for  volume,  than  water,  a 
sinker  is  attached,  whose  weight  in  water  =  x,  and  rel.  dens.  =  cL 


96  APPENDIX. 

The  relative  density  of  the  substance  lighter  than  water  is  then 
expressed  by  the  formula 


where  W"  =  weight  of  sinker  and  solid  in  water. 
If  the  relative  density  be  required  at  t°  C.,  then 

Eel.  dens.  =      ^      x  rel.  dens,  of  water  at  t°  C. 

The  relative  density  of  a  liquid  is  commonly  found  by 
1st.  The  specific  gravity  flask  method. 

Let  x  =  weight  of  flask  empty,  W=  weight  of  flask  filled 
with  water  at  t°  C.,  W'=  weight  of  flask  filled  with  liquid  under 
examination;  then 

-nri  _  x 

Eel.  dens.  =  —  -  x  rel.  dens,  of  water  at  t°  C. 
W—  x 

2d.  By  weighing  a  solid  of  constant  volume  in  water  and 
then  in  the  liquid. 

Let  cc  =  weight  of  solid  in  air,  W=  weight  in  water  at  £°C., 
W'=  weight  in  liquid;  then 

g.    _    TTTf 

Eel.  dens.  =  -  x  rel.  dens,  of  water  at  t°  C. 
x-W 


Density  of  Gases  and  Vapors. 

The  specific  gravity,  or  relative  density  of  a  gas  or  vapor, 
as  has  been  shown  on  pp.  18  and  19,  is  the  ratio  of  its  mass  to 
the  mass  of  an  equal  volume  of  hydrogen,  measured  at  the 
same  temperature  and  pressure.  By  the  density  of  gas  or 
vapor,  we  mean  the  relative  density. 


APPENDIX.  97 

One  liter  of  hydrogen  gas  at  0°  C.,  and  760  mm.  barometric 
pressure  at  the  sea-level  weighs  .0896  gram. 

The  relative  density  of  a  gas  is  determined  by  weighing  a 
known  volume  of  the  gas,  and  comparing  it  with  the  weight  of 
an  equal  volume  of  hydrogen  under  like  conditions  of  temper- 
ature and  pressure  (p.  21).  The  effusion  method  may  also 
be  used.  "The  rel.  dens,  varies  directly  as  the  square  of  the 
time  of  effusion  of  equal  volumes."  Graham's  Law  (see  p.  25) 
may  also  be  applied. 

EXAMPLES. 

1.  Calculate  the  relative  density  of  a  solid  from  the  follow- 
ing data :  — 

Weight  of  substance  in  air,  2.4554  grams. 
Weight  of  substance  in  water,  2.0778  grams. 

2.  Determine  the  relative  density  of  wood  from  the  follow- 
ing data : — 

Weight  of  wood  in  air,  4  grams. 
Weight  of  silver  sinker  in  air,  10  grains. 
Weight  of  wood  and  sinker  under  water,  8.5  grams. 
Relative  density  of  silver  =  10.5. 

3.  A  solid  weighs  in  vacuo  100  grams,  in  water  85  grams, 
and  in  another  liquid  88  grams.     What  is  the  relative  density 
of  this  liquid? 


. 

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with  the  constellations.  47  pages.  Paper.  15  cts. 

Guides  for  Science  Teaching.  Teachers'  aids  in  the  instruction  of  Natural  History 
classes  in  the  lower  grades. 

I.     Hyatt's  About  Pebbles.     26  pages.     Paper.     10  cts. 
II.     Goodale's  A  Few  Common  Plants.     61  pages.     Paper.     20  cts. 

III.  Hyatt's  Commercial  and  other  Sponges.    Illustrated.    43  pages.  Paper.   20  cts. 

IV.  Agassiz's  First  Lessons  in  Natural  History.     Illustrated.     64  pages.     Paper. 

25  cts. 

V.     Hyatt's  Corals  and  Echinoderms.     Illustrated.     32  pages.    Paper.     30  cts. 
VI.     Hyatt's  Mollusca.     Illustrated.     65  pages.     Paper.     30  cts. 
VII.      Hyatt's  Worms  and  Crustacea.     Illustrated.     68  pages.     Paper.     30  cts. 
VIII.     B  yatt's  Insecta.     Illustrated.     324  pages.     Cloth.     #1.25. 
XII.    Crosby's  Common  Minerals  and  Rocks.     Illustrated.    200  pages.     Paper,  49^ 

cts.     Cloth,  60  cts. 

XIII      .niichard's  First  Lessons  in  Minerals.     50  pages.     Paper.     10  cts. 
XIV.     Bowditch's  Physiology.     58  pages.     Paper.     20  cts. 
XV      Clapp's  36  Observation  Lessons  in  Minerals.    80  pages.     Paper.     30  cts. 
XVI,     Phenix's  Lessons  in  Chemistry.     In  press. 
Pupils'  Note-Book  to  accompany  No.  15.     10  cts. 

Rice's  Science  Teaching  in  the  School.  With  a  course  of  instruction  in  science 
for  the  lower  grades.  46  pages.  Paper.  25  cts. 

Ricks'?  Natural  History  Object  Lessons.  Supplies  information  on  plants  and 
their  products,  on  animals  and  their  uses,  and  gives  specimen  lessons.  Fully  illustrated. 
332  pages.  $1.50. 

Ricks's  Object  Lessons  and  How  to  Give  them. 

Volume  I.     Gives  lessons  for  primary  grades.     200  pages.     90  cts. 

Volume  II.  Gives  lessons  tor  grammar  and  intermediate  grades.  212  pages.     90  cts. 

Shaler's  First  Book  in  Geology.  For  high  school,  or  highest  class  in  grammar  school. 
"*72  pages.  Illustrated.  $1.00. 

Shaler's  Teacher's  Methods  in  Geology.  An  aid  to  the  teacher  of  Geology. 
74  pages.  Paper.  25  cts. 

Smith's  Studies  in  Nature.  A  combination  of  ratural  history  lessons  and  language 
work.  48  pages.  Paper.  15  cts. 

Sent  by  mail  postpaid  on  receipt  of  price.     See  also  our  list  of  books  in  Science. 


D.    C.    HEATH   &   CO.,    PUBLISHERS, 

BOSTON.         NEW  YORK.        CHICAGO. 


MATHEMATICS. 


Bowser's  Academic  Algebra.  A  complete  treatise  through  the  progressions,  includ 
ing  Permutations,  Combinations,  and  the  Binomial  Theorem.  Half  leather.  $1.15. 

Bowser's  College  Algebra.  A  complete  treatise  for  colleges  and  scientific  schools. 
Half  leather.  $1.65. 

Bowser's  Plane  and  Solid  Geometry.  Combines  the  excellences  of  Euclid  with 
those  of  the  best  modern  writers.  Half  leather.  $1.35. 

Bowser's  Plane  Geometry.    Half  leather.    85  cts. 

Bowser's  Elements  of  Plane  and  Spherical  Trigonometry.     A  brief  course 

prepared  especially  for  High  Schools  and  Academies.     Half  leather.     $1.00. 

Bowser's  Treatise  on  Plane  and  Spherical  Trigonometry.    An  advanced 

work  which  covers  the  entire  course  in  higher  institutions.     Half  leather.    $1.65. 

Hanus's  Geometry  in  the  Grammar  Schools.    An  essay,  together  with  illustrative 

class  exercises  and  an  outline  of  the  work  for  the  last  three  years  of  the  grammar  school. 
52  pages.     25  cts. 

Hopkin's  Plane  Geometry.     On  the  heuristic  plan.     Half  leather.     85  cts. 

Hunt's  Concrete  Geometry  for  Grammar  Schools.  The  definitions  and  ele- 
mentary concepts  are  to  be  taught  concretely,  by  much  measuring,  by  the  making  of 
models  and  diagrams  by  the  pupil,  as  suggested  by  the  text  or  by  his  own  invention, 
loo  pages.  Boards.  30  cts.  » 

Waldo's  Descriptive  Geometry.  A  large  number  of  problems  systematically  ar- 
ranged and  with  suggestions.  90  cts. 

The  New  Arithmetic.  By  300  teachers.  Little  theory  and  much  practice.  Also  aa 
excellent  review  book.  230  pages.  75  cts. 

For  Arithmetics  and  other  elementary  work  see  our  list  of  books  in  Number. 


D.   C.   HEATH    &   CO.,  PUBLISHERS, 

BOSTON.        NEW  YORK.        CHICAGO. 


ENGLISH  LANG  UA  GE. 


Hyde's  Lessons  in  English,  Book  I.  For  the  lower  grades.  Contains  exercises 
for  reproduction,  picture  lessons,  letter  writing,  uses  of  parts  of  speech,  etc.  40  cts. 

Hyde's  Lessons  in  English,  BOOk  II.  For  Grammar  schools.  Has  enough  tech- 
nical  grammar  for  correct  use  of  language.  60  cts. 

Hyde's  Lessons  in  English,  Book  II  with  Supplement.     Has,  in  addition 

to  the  above,  118  pages  of  technical  grammar.     70  cts. 
Supplement  bound  alone,  35  cts. 

Hyde's  Advanced  Lessons  in  English.  For  advanced  classes  in  grammar  schools 
and  high  schools.  60  cts. 

Hyde's  Lessons  in  English,  Book  II  with  Advanced  Lessons.     The  Ad- 

vanced  Lessons  and  Book  II  bound  together.     80  cts. 

Hyde's  Derivation  of  Words.     15  cts. 

Mathews's  Outline  of  English  Grammar,  with  Selections  for  Practice. 

The  application  of  principles  is  made  through  composition  of  original  sentences.     80  cts. 
Buckbee's   Primary  Word  BOOk.      Embraces  thorough  drills  in  articulation  and  in 
the  primary  difficulties  of  spelling  and  sound.     30  cts. 

Sever's  Progressive  Speller.  For  use  in  advanced  primary,  intermediate,  and  gram- 
mar  grades.  Gives  spelling,  pronunciation,  definition,  and  use  of  words.  30  cts. 

Badlam's  Suggestive  Lessons  in  Language.    Being  Part  I  and  Appendix  of 

Suggestive  Lessons  in  Language  and  Reading.     50  cts. 

Smith's  Studies  in  Nature,  and  Language  Lessons.    A  combination  of  object 

lessons  with  language  work.     50  cts.     Part  I  bound  separately,  25  cts. 

Meiklejohn's  English  Language.  Treats  salient  features  with  a  master's  skill  and 
with  the  utmost  clearness  and  simplicity.  $1.30. 

Meiklejohn's  English  Grammar.  Also  composition,  versification,  paraphrasing,  etc. 
For  high  schools  and  colleges.  90  cts. 

Meiklejohn's  History  of  the  English  Language.  78  pages.  Part  III  of  Eng- 
lish Language  above,  35  cts. 

Williams' s  Composition  and  Rhetoric  by  Practice.    For  high  school  and  col- 

lege.     Combines  the  smallest  amount  of  theory  with  an  abundance  of  practice.    Revised 
edition.     $1.00. 

Strang's  Exercises  in  English.  Examples  in  Syntax,  Accidence,  and  Style  for 
criticism  and  correction.  50  cts. 

HuffCUtt's  English  in  the  Preparatory  School.  Presents  as  practically  as  pos- 
sible some  of  the  advanced  methods  of  teaching  English  grammar  and  composition  in  the 
secondary  schools.  25  cts. 

Woodward's  Study  Of  English.  Discusses  English  teaching  from  primary  school  to 
high  collegiate  work.  25  cts. 

Gemmg's  Study  Of  Rhetoric.  Shows  the  most  practical  discipline  of  students  for  tin 
making  of  literature.  25  cts. 

GC2dchild'S  BOOk  Of    Stops.     Punctuation  in  Verse.     Illustrated.     10  cts. 
See  also  our  list  of  books  for  the  study  of  English  Literature. 


D.    C.    HEATH    &    CO.,    PUBLISHERS, 

BOSTON.        NEW  YORK.        CHICAGO. 


ENGLISH  LITERATURE. 

Hawthorne  and  Lemmon's  American  Literature.    A  manual  for  high  schools 

and  academies.     £1.25. 

Meiklejohn's  History  of  English  Language  and  Literature.    For  high  schools 

and  colleges.     A  compact   and  reliable  statement   of   the   essentials  ;  also  included  in 
Meiklejohn's  English  Language  (see  under  English  Language).     90  cts. 

Meiklejohn's  History  of  English  Literature.     n6  pages.     Part  iv  of  English 

Literature,  above.     45  cts. 

Hodgkins'  Studies  in  English  Literature.  Gives  full  lists  of  aids  for  laboratory 
method.  Scott,  Lamb,  Wordsworth,  Coleridge,  Byron,  Shelley,  Keats,  Macaulay» 
Dickens,  Thackeray,  Robert  Browning,  Mrs.  Browning,  Carlyle,  George  Eliot,  Tenny- 
son, Rossetti,  Arnold,  Ruskin,  Irving,  Bryant,  Hawthorne,  Longfellow,  Emerson, 
Whittier,  Holmes,  and  Lowell.  A  separate  pamphlet  on  each  author.  Price  5  cts.  each, 
or  per  hundred,  $3.00;  complete  in  cloth  (adjustable  file  cover,  #1.50).  $1.00. 

Scudder's  Shelley's  Prometheus  Unbound.      With  introduction  and  copious 

notes.     70  cts. 

George's  Wordsworth's  Prelude.  Annotated  for  high  school  and  college.  Never 
before  published  alone.  80  cts. 

George's  Selections  from  Wordsworth.  168  poems  chosen  with  a  view  to  illustrate 
the  growth  of  the  poet's  mind  and  art.  $1.00. 

George's  Wordsworth's  Prefaces  and  Essays  on  Poetry.    Contains  the  best  of 

Wordsworth's  prose.    60  cts. 
George's  Webster's  Speeches.      Nine  select  speeches  with  notes.     $1.50. 

George's  Burke' s  American  Orations.    Cloth.    65  cts. 

George's  Syllabus  of  English  Literature  and  History.     Shows  in  parallel 

columns,  the  progress  of  History  and  Literature.     20  cts. 

CorSOn'S  Introduction  tO  Browning.  A  guide  to  the  study  of  Browning's  Poetry. 
Also  has  33  poems  with  notes.  $1.50. 

Corson's  Introduction  to  the  Study  of  Shakespeare.    A  critical  study  of 

Shakespeare's  art,  with  examination  questions.     $1.50. 

Corson's  Introduction  to  the  Study  of  Milton,     in  press. 
Corson's  Introduction  to  the  Study  of  Chaucer,     in  press. 

COOk's  Judith.  The  Old  English  epic  poem,  with  introduction,  translation,  glossary  and 
fac-simile  page.  $1.60.  Students'  edition  without  translation.  35  cts. 

COOk's  The  Bible  and  English  Prose  Style.  Approaches  the  study  of  the  Bible 
from  the  literary  side.  60  cts. 

Simonds'  Sir  Thomas  Wyatt  and  his  Poems.     168  pages.    With  biography,  and 

critical  analysis  of  his  poems.     75  cts. 
Hall's  BeOWUlf.      A  metrical  translation.     $1.00.     Students'  edition.     35  cts. 

Norton's  Heart  Of  Oak  BOOkS.  A  series  of  five  volumes  giving  selections  from  the 
choicest  English  literature. 

Phillips's  History  and  Literature  in  Grammar  Grades.    An  essay  showing  the 

intimate  relation  of  the  two  subjects.     15  cts. 

See  also  our  list  of  books  for  the  study  of  the  English  Language. 


D.  C.  HEATH  &  CO.,  PUBLISHERS. 

BOSTON.        NEW  YORK.        CHICAGO. 


HISTORY. 


Sheldon's  United  States  History.  For  grammar  schools.  Follows  the  "  seminary  " 
or  laboratory  plan.  "  By  it  the  child  is  not  robbed  of  the  right  to  do  his  own  think- 
ing." Half  leather.  $1.25. 

Teacher's  Manual  to  Sheldon's  United  States  History.   A  key  to  the  above 

system.     60  cts. 

Sheldon's  General  History.  For  high  school  and  college.  The  only  general  history 
following  the  "seminary"  or  laboratory  plan  now  advocated  by  leading  teachers. 
Half  leather.  $1.75. 

Sheldon's  Greek  and  Roman  History.  Contains  the  first  250  pages  of  the  above 
book.  $1.00. 

Teacher's  Manual  tO  Sheldon's  History.  Puts  into  the  instructor's  hand  the  key 
to  the  above  system.  85  cts. 

Sheldon's  Aids  to  the  Teaching  of  General  History.   Gives  also  list  of  most 

essential  books  for  a  reference  library.     10  cts. 

Thomas's  History  Of  the  United  States.     For  schools,  academies,  and  the  general 
reader.     A  narrative  history  with  copious  re" 
illustrated.     532  pages.     Half  leather.     $1.25. 


reader.     A  narrative  history  with  copious  references  to  sources  and  authorities.     Fully 
"ilflea  • 


Shumway's  A  Day  in  Ancient  Rome.  With  59  illustrations.  Should  find  a  place 
as  a  supplementary  reader  in  every  high-school  class  studying  Cicero,  Horace,  Taci- 
tus, etc.  75  cts. 

Old  South  Leaflets.  Reproductions  of  important  political  and  historical  papers,  ac- 
companied by  useful  notes.  Each,  5  cts.  and  6  cts.  For  titles  see  separate  lists.  Per 
hundred,  $3.00. 

Allen's  History  Topics.  Covers  Ancient,  Modern,  and  American  history,  and  gives  an 
excellent  list  of  books  of  reference.  121  pages.  Paper.  30  cts. 

Fisher's  Select  Bibliography  of  Ecclesiastical  History.   An  annotated  list  of 

the  most  essential  books  for  a  theological  student's  library.     15  cts. 

Hall's  Method  Of  Teaching  History.  "Its  excellence  and  helpfulness  ought  to 
secure  it  many  readers." — The  Nation.  $1.50. 

Phillips'  History  and  Literature  in  Grammar  Grades.    A  paper  read  before  the 

Department  of  Superintendence,  at  Brooklyn,  N.Y.     Paper.     15  cts. 


See  also  our  list  of  Old  South  Leaflets. 


D.    C.   HEATH    &   CO.,  PUBLISHERS, 

BOSTON.         NEW  YORK.        CHICAGO. 


CIVICS,  ECONOMICS,  AND  SOCIOLOGY. 


Boutwell's  The  Constitution  of  the  United  States  at  the  End  of  the  First 

Century.  Contains  the  Organic  Laws  of  the  United  States,  with  references  to  the 
decisions  of  the  Supreme  Court  which  elucidate  the  text,  and  an  historical  chapter  re- 
viewing the  steps  which  led  to  the  adoption  of  these  Organic  Laws.  In  press. 

Dole's   The  American  Citizen.     Designed  as  a  text-book  in  Civics  and  morals  for  the 

higher  grades  of  the  grammar  school  as  well  as  for  the  high  school  and  academy.  Con- 
tains Constitution  of  United  States,  with  analysis.  336  pages.  $1.00. 

Special  editions  are  made  for  Illinois,  Indiana,  Ohio,  Missouri,  Nebraska,  No.  Dakota, 
So.  Dakota,  Wisconsin,  Minnesota,  Kansas,  Texas. 

Goodale's  Questions  to  Accompany  Dole's  The  American  Citizen.  Con- 
tains, beside  questions  on  the  text,  suggestive  questions  and  questions  for  class  debate. 
87  pages.  Paper.  25  cts. 

Gide's  Principles  Of  Political  Economy.  Translated  from  the  French  by  Dr. 
Jacobsen  of  London,  with  introduction  by  Prof.  James  Bonar  of  Oxford.  598 
pages.  $2.00. 

Henderson's  Introduction  to   the  Study  of   Dependent,  Defective,  and 

Delinquent  Classes.  Adapted  for  use  as  a  text-book,  for  personal  study,  for 
teachers'  and  ministers' institutes,  and  for  clubs  of  public-spirited  men  and  women  engaged 
in  considering  some  of  the  gravest  problems  of  society.  287  pages.  $1.50. 

Hodgin's  Indiana  and  the  Nation.  Contains  the  Civil  Government  of  the  State, 
as  well  as  that  of  the  United  States,  with  questions.  198  pages.  70  cts. 

Lawrence's  Guide  to  International  Law.    A  brief  outline  of  the  principles  and 

practices  of  International  Law.     In  press. 

Wenzel's  Comparative  View  Of  Governments.  Gives  in  parallel  columns  com- 
parisons of  the  governments  of  the  United  States,  England,  France,  and  Germany.  26 
pages.  Paper.  22  cts. 

Wilson's  The  State.  Elements  of  Historical  and  Practical  Politics.  A  text-book  on 
the  organization  and  functions  of  government  for  high  schools  and  colleges.  720  pages. 

$2.00. 

Wilson's  United  States  Government.  For  grammar  and  high  schools.  140  pages. 
60  cts. 

Woodburn  and  Hodgin's  The  American  Commonwealth.     Contains  several 

orations  from  Webster  and  Burke,  with  analyses,  historical  and  explanatory  notes,  and 
studies  of  the  men  and  periods.  586  pages.  $1.50. 

Sent  by  mail,  post  paid  on  receipt  of  prices.  See  also  our  list  of  books  in  History. 


D.    C.    HEATH    &    CO.,    PUBLISHERS, 

BOSTON.        NEW  YORK.        CHICAGO. 


EDUCATION. 

Compayre''s  History  Of  Pedagogy.  "  The  best  and  most  comprehensive  history  of 
Education  in  English."  —  Dr.  G.  S.  HALL.  $1.75. 

CompayrS's  Lectures  On  Teaching.  "  The  best  book  in  existence  on  thr  theory  an<J 
practice  of  education."  —  Supt.  MACALISTER,  Philadelphia.  $1.75. 

Compayr6'8  Psychology  Applied  tO  Education.  A  clear  and  concise  statement 
of  doctrine  aud  application  on  the  science  and  art  of  teaching.  90  cts. 

De  Garmo's  Essentials  Of  Method.  A  practical  exposition  of  methods  with  illustra- 
tive outlines  of  common  school  studies.  65  cts. 

De   GarmO's   Lindner's   Psychology.      The  best   Manual  ever  prepared  from  the 

Herbartian  standpoint.     $1.00. 
Gill's  Systems  Of  Education.      "  It  treats  ably  of  the  Lancaster  and  Bell  movement 

in  education,  —  a  very  important  phase."  —  Dr.  W.  T.  HARRIS.     $1.25. 

Hall's  Bibliography  Of  Pedagogical  Literature.      Covers  every  department  of 

education.     Interleaved,  *$2.oo.     $1.50. 
Herford'S  Student's  Froebel.      The  purpose  of  this  little  book  is  to  give  young  people 

preparing  to  teach  a  brief  yet  full  account  of  Froebel's  Theory  of  Education.     75  cts. 

Malleson's  Early  Training  of  Children.    "The  best  book  for  mothers  I  ever 

read."  —  ELIZABETH  P.  PEABODY.     75  cts. 

Marwedel's    Conscious   Motherhood.     The  unfolding  of  the  child's   mind  in  the 

cradle,  nursery  and  Kindergarten.     $2.00. 
Newsholme'S  School  Hygiene.     Already  in  use  in  the  leading  training  colleges  in 

England.     75  cts. 

Peabody's  Home,  Kindergarten,  and  Primary  School.  "The  best  book  out- 
side of  the  Bible  that  I  ever  read."  — A  LEADING  TEACHER.  JSi.oo. 

PestalOZZi'S  Leonard  and  Gertrude.  "If  we  except  'Emile'  only,  no  more  im- 
portant educational  book  has  appeared  for  a  century  and  a  half  than  '  Leonard  and  Ger- 
trude.' "  —  The  Nation.  90  cts. 

RadestOCk'S  Habit  in  Education.  "  It  will  prove  a  rare  '  find '  to  teachers  who  are 
seeking  to  ground  themselves  in  the  philosophy  of  their  art."  —  E.  H.  RUSSELL,  Worces- 
ter Normal  School.  75  cts. 

Richter's  Levana  ;  or,  The  Doctrine  of  Education.     "A  spirited  and  scholarly 

•book."  — Prof.  W.  H.  PAYNE.    $1.40. 

Rosmini'S  Method  in  Education.  "The  most  important  pedagogical  work  ever 
written."  — THOMAS  DAVIDSON.  $1.50. 

ROUSSeau'S  Emile.      "  Perhaps  the  most  influential  book  ever  written  on  the  subject  of 

Education."  —  R.  H.  QUICK.     90  cts. 
Methods  Of  Teaching  Modern  Languages.      Papers  on  the  value  and  on  methods 

of  teaching  German  and  French,  by  prominent  instructors.     90  cts. 

Sanford's  Laboratory  Course  in  Physiological  Psychology.     The  course 

includes  experiments  upon  the  Dermal  Senses,  Static  and  Kinaesthetic  Senses,  Taste, 
Smell,  Hearing,  Vision,  Psychophysic.     In  Press. 

Lange's  Apperception :  A  monograph  on  Psychology  and  Pedagogy.  Trans- 
lated by  the  members  of  the  Herbart  Club,  under  the  direction  of  President  Charles 
DeGarmo,  of  Swarthmore  College.  $>i.oo. 

Herbart' S  Science  Of  Education.  Translated  by  Mr.  and  Mrs.  Felken  with  a  pref- 
ace by  Oscar  Browning.  $1.00. 

Tracy's   Psychology   Of  Childhood.    This  is  the  first  general  treatise  covering  in  a 
scientific  manner  the  whole  field  of  child  psychology.     Octavo.     Paper.     75  cts. 
Sent  by  mail,  postpaid,  on  receipt  of  price. 

D.    C.    HEATH    &    CO.,    PUBLISHERS, 

BOSTON.        NEW  YORK.        CHICAGO. 


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